all digits other than zero are significant
zeroes between non-zero digits are significant
final zeros to the right of a decimal point are significant
in numbers smaller than one, zeroes directly to the left or right of the decimal point are not significant
zeroes at the end of a measurement may or may not be significant. To avoid uncertainty follow this rule:

B. ROUND OFF TO THE NUMBER OF SIG. FIG. SHOWN AND PLACE IN SCIENTIFIC NOTATION

C. Rules for Adding and Subtracting

When adding or subtracting, the numbers should have the same number of decimal places as the quantity having the least number of decimal places.

D. Rules for Multiplication and Division

The number of significant figures in a product or a quotient obtained from measured quantities is the same as the number of significant figures in the quantity having the smaller number of significant figures.

Home

Notes- Overview of course

I. Integrated Physics and Chemistry

a. Study of non-living things dealing mainly with chemistry and physics

b. Chemistry- The Study of the structure and composition of substances and the changes they will undergo

c. Physics- Study of matter and energy and their interactions; concerned with the description and understanding of nature

II. Scientific method

a. Helps scientists answer questions about the physical universe

1. Observations

i. made when they note and record facts

2. Hypothesis

i. Based on observations; tentative explanations which provides direction for future investigation

3. Experiments

i. One or more conditions are controlled; can be repeated; manipulated variable-is something that is intentionally changed

4. Theories

i. A former hypothesis that has been tested with repeated experiments; reasonable explanation from observed events that are related; involves an imaginary model

ii. Model- description of something unfamiliar in terms of something familiar

5. Laws

i. A generalized statement that describes behavior; describes a natural event stated in words; could be stated in a mathematical equation

Home

Notes- Metric System

I. International System (SI)

a. Root words

1. meter-length

2. liter (or cubic meter)- volume

3. kilogram (gram for metric system)- mass

b. Prefixes

1. Kilo- 10³

2. Hecto-10²

3. Deka-10¹

4. Root-10⁰

5. deci-10^-1

6. centi-10^-2

7. milli-10^-3

II. Conversions

English	Metric
1 in	2.54 cm
39.4 in	1 meter
1 lb	454 g
2.2 lb	1 kg
1.06 qt	1 L
1 qt	946 mL
1 mi	1.62 km
.621 mi	1 km
1 oz	29.6 mL

Kilo	Hecto	Deka	Root	Deci	Centi	Milli
Km	Hm	Dkm	M	Dm	Cm	Mm
Kl	Hl	Dkl	L	Dl	Cl	Ml
Kg	Hg	Dg	G	Dg	Cg	Mg

ALWAYS MEASURE TO THE TENTHS PLACE

Area- amount of surface space contained within a given set of lines

A=LxW

Units are squared

Volume- amount of space occupied by some form of matter

Regular objects- LxWxH

Irregular objects- Water Displacement

1 cm³=1 mL

Home

Notes- Conversion Factors and Dimensional Analysis

I. Conversion Factors

a. Relation between 2 equivalent measurements; can be multiplied by a measurement without changing its true value

II. Dimensional Analysis

a. Helps you tell whether you are using the correct equation in solving a problem (Showing work)

Home

Notes- Density and Specific Gravity

I. Density

a. Mass of a substance per unit of volume

b. D=M/V

c. Unit- g/cm³ or g/mL or kg/m³

d. Density or water is 1 g/mL or 1 g/cm³ at 25°C

II. Specific Gravity

a. Ratio of density of the substance to the density of water

b. SpG= density of a substance/ density of water

c. No Units

d. Determines whether something will sink or float

e. If the SpG is greater than 1, then it will sink

f. If the SpG is less than 1, then it will float

Home

Notes- Properties n’ Change

I. States of matter

a. Particle arrangement

i. Closely spaced (solids)

1. crystalline solid - looks like crystal

2. amorphous solid

a. slow moving solid

b. change in temperature

c. candles

d. glass

ii. Liquid - particles farther apart

iii. Gas

1. fill entire container

2. particles very far apart

iv. Plasma

1. exists at extremely high temperatures

2. stars

3. molecules dissociate

4. atoms become ionized

v. Bose - Einstein model

1. Super cool gas

2. acts different

II. Properties

a. Physical properties

i. Properties that can be observed with our senses, and/or instruments

1. Size

2. Shape

3. Color

4. Luster

5. Hardness

6. Malleability

7. Conductivity

8. Ductility

9. Boiling point

10. Freezing point

11. Taste

12. Smell

13. Texture

14. Density

b. Chemical properties

i. Rust

ii. Burn

iii. React with acids

iv. React with other substances

III. Changes

a. Physical changes affect appearance

i. Changes

1. size

2. shape

3. state

ii. State Change

1. particles don’t change, they rearrange

2. solid à gas – sublimation

3. liquidà gas – vaporization

a. evaporation-occurs at top of liquid at any time

b. boiling- occurs throughout when boiling point is reached

4. gas à solid - deposition

5. liquid à solid - freezing

6. solid à liquid - melting

7. gasà liquid – condensation

b. Law of conservation of matter: Matter can not be created nor destroyed

c. Chemical changes

i. Does at least one of the following

1. forms a precipitate

2. production of a gas

3. permanent color change

4. heat and light are produced

Home

Notes- Solutions

I. Components of a solution

a. “To Loosen” (Latin)

b. Examples

i. Vinegar- 3% Acetic Acid and Water

ii. Ocean

iii. Blood Serum

c. Solid Solutions

i. Alloys

1. mixture of metals

2. has higher strength

3. greater resistance to corrosion

4. higher melting point

5. brass- copper and zinc

6. 14 karat gold- gold, silver, copper

d. Gas

i. Air

e. Solution- Type of mixture

f. Mixtures can be physically separated

g. Solute- stuff being dissolved

h. Solvent- substance that does the dissolving

i. Solvent is always greater in amount than the solute

II. Type of solutions

a. Solutions are homogeneous mixtures

i. Can’t see the different parts

b. Heterogeneous Mixture

i. Can see different parts

ii. Can be separated by filtering

Homogeneous	Colloids	Heterogeneous
Solutions- mixtures where particles are too small to be seen	Appears to be homogenous, but is classified as a heterogeneous	Suspensions- made up of finely divided solid matter in a liquid which the solid will settle out on standing
Air	Don’t separate on standing	Muddy water, oil-based paints, Pepto-Bismol
Saltwater	Particles 1à 100 nanometers	Will scatter light
Alloys	Will scatter light	100 nanometers or larger
Particles range from .1 à 1 nanometer	Milk, toothpaste, butter, mayonnaise, blood

III. Rate of solutions

a. Crushing

i. Increases surface area

ii. Hydration: when water is solvent

iii. Diffusion- spreading of solute particles from a concentrated area to a less concentrated areas

b. Stirring

i. Adds energy

ii. Speeds up dissolving rate

c. Heating

i. Raises temperature to increase energy

d. Rate of Solution- how fast something dissolves

IV. Unsaturated, Saturated, and Supersaturated

a. Unsaturated

i. More solute can be dissolved

ii. Dilute- small amount of solute

iii. Concentrated- large amount of solute

b. Saturated

i. Cannot dissolve any more solute

ii. Precipitate- solid that separates from solution

iii. Crystal remains dissolved when solution cools

c. Supersaturated

i. More solute than expected at a given temperature is dissolved

ii. Unstable

iii. Agitate might fall out

V. Factors affecting solubility

a. Most things dissolve quicker in something hot

b. All substances have a solubility factor

c. Solubility- The maximum amount of a solute that will dissolve in a certain amount of solvent at a given temperature.

d. To increase solubility in a liquid, heat it

e. Insoluble- will not dissolve in water

f. Amount of gas dissolved decreases as temperature increases

g. Extra energy causes gas to escape

h. Increase pressure, increase solubility of a gas in a liquid.

VI. Properties of a solute

a. Water does not conduct electricity

b. Solution can conduct electricity if it has ions in it.

c. Solutes raise the boiling point of a liquid solvent

d. Pure water freezes quicker than a water solution

e. Solutes may release heat when dissolving

f. Exothermic- releases heat energy

g. Endothermic- absorbs heat energy

VII. Percent By Weight Formula

a. Mass of solute/mass of solution x100

Home

Notes- Elements, Atoms, Compounds

A. Elements

Definition- pure substance that cannot be broken down by physical and chemical changes.

a. Nuclear energy can produce new substance not there before

Dalton’s Atomic Theory

a. In the early 1800’s John Dalton made a theory about the atom that stated the following:

1. Matter is composed of tiny particles called atoms that are indivisible

a. False, because of subatomic particles: protons, neutrons, and electrons

2. Atoms of the same element are alike

a. True, because number of protons in the atom do not change

b. False, because the number of neutrons may vary in atoms of the same elements. All elements have isotopes.

1. 3 isotopes of hydrogen

a. H 1pt 1e- 0_n – Protium

b. H 1pt 1e- 1_n -Deuterium

c. H 1pt 1e- 2_n -Tritium

3. Atoms of different elements are different

a. True, because number of proton changes from element to element, which determines what element it is.

4. During chemical changes, the atoms unite in a definite proportion to form a new substance

a. True, because atoms of a compound always combine in a certain ratio. Ex: H₂0, CO₂, and NaCl.

5. During chemical changes, the atoms themselves do not change

a. True - reactions have same atoms but rearranged

Chemical Symbols-a short hand method for writing elements.

a. Jakob Berzelius - devised system for writing chemical symbols

1. First letter is always Capital

2. Second letter is always lower case

3. 28 element names and symbols

B. Compounds

Definition - Compounds are pure substances that are made of two or more elements chemically combined
Bonding

a. Covalent - means atoms share electrons. When atoms share electrons called molecules. Covalent are only molecules. Ex. Water.

b. Ionic - atoms transfers electrons when atoms do this they form ions. Ex. Table salt.

Diatomic Elements

a. Definition - Two atoms of the same element join covalently to have all of the properties of that element; these diatomic elements always exist in pairs in nature.

b. Examples:

1. Nitrogen-N₂

2. Oxygen-O₂

3. Fluorine-F₂

4. Chlorine-Cl₂

5. Bromine-Br₂

6. Iodine-I₂

7. Astatine-At₂

8. Hydrogen-H₂

Home

Notes- Atomic Structure

A. Historical Model of the Atom

Scientists had used indirect methods to describe the atom
Timeline in Atomic Model

a. Greeks believe that if matter was divided until eventually it could not be further divided, then that was the atom. Atom means indivisible in Greek.

b. 1800 - John Dalton’s models the atom as solid spheres.

c. 1897- J.J. Thompson discovers the electron and includes it in revised version of atom model. Described the atom as a “ball of electricity.”

d. 1909 - Robert Millikan determines the mass of an electron; measured at 9.109 X 10^-28 g

e. 1911 - Lord Rutherford discovers dense positively charged mass in center of atom called nucleus.

1. Fired positive particle, radium alpha particle, into gold foil surrounded by photographic film. Results showed that most of particles went through but some were deflected back.

2. Used Thomson’s model combined with the nucleus.

3. Disproved Thomson’s model and said the following:

a. Atoms contains nucleus and electrons

b. Nucleus is tiny dense positively charged

c. Atom is mostly empty space

d. Credited with finding the nucleus; not protons or neutrons

e. Mass of proton is 1.673 X 10^-24 g

f. 1913 - Neils Bohr - proposes electrons orbit around nucleus. Method still used to diagram atoms

g. 1932 - James Chadwick - discovered the neutron

1. Mass is 1.675 X 10^-24 g

h. Electron Cloud Model

1. Atoms are organized around a dense nucleus with electrons traveling in energy level

B. Protons, Electrons, and Neutrons

Nucleus are made of two types of subatomic particles

a. Protons which carry a positive charge

b. Neutrons which carry no charge

Electrons are subatomic particles with a negative charge

a. Most of mass is made up of protons and neutrons

b. Electrons have almost no mass

Atoms as a whole is neutral (no charge)

a. Number of protons equal number of electrons

b. Exception is ions

C. Atomic Mass and Mass Number

AMU- atomic mass unit

a. 1 proton= 1 AMU

b. 1 neutron= 1 AMU

Mass number is number of protons and neutrons for one of the isotopes of that element. Includes either most stable or most common.
Atomic mass - the average mass of all of the isotopes for that element.

a. Francis William Aston used mass spectrometer and was able to determine the atomic masses to the fourth decimal place.

D. Organization of the Periodic Table

Elements in each column have same number of valence electron
The Roman numeral (I-VII) indicate number of valence electrons (in our period table it is all labeled with A)
Elements in each column have similar physical and chemical properties.
First published periodic table by Dmitri Mendeelev in 1864

a. organized by atomic masses

Henry Mosely put the table by atomic numbers
IUPAC- International Union of Pure and Applied Chemistry

a. Consists of chemists from over 80 different countries

Columns are groups or families that are 18 in number (vertical)
Periods are rows that are seven in number (horizontal)
Alkaline Metals

a. Found in Group 1 in the periodic table

b. Soft, silver white metals

c. React readily with other elements

d. Never found by themselves

e. Most reactive metal: Includes Li, Na, K, Rb, Cs, and Fr

Alakaline Earth Metals

a. Group 2 in Periodic Table

b. 2 valence electrons

c. React readily with other elements

d. Rarely found free in nature: Includes Be, Mg, Ca, Sr, Ba, and Ra

Transition Metals (B’s on our periodic table)

a. Group 3-12

b. They can have different number of valence electrons

c. High melting points, with W (tungsten) having the highest of any element

d. High luster; can be polished

Boron Family

a. Group 13; have 3 valence electrons

b. Boron (B) is a semi-metal; all others are metals

Carbon Family

a. Group 14 (or 4a in our periodic table); have 4 valence electrons

b. Carbon is the only non-metal; Silicon (Si) and Germanium (Ge) are semi-metals; Tin (Sn) / Lead (Pb) are only metals

Nitrogen Family

a. Group 15 (or 5a in our periodic table); have 5 valence electrons

b. Nitrogen (N) and Phosphorus (P) are the only non-metals; Arsenic (As) and Antimony (Sb) are semi-metals; Bismuth (Bi) is only metal

Oxygen Family

a. Group 16 (or 6a in our periodic table); have 6 valence electrons

b. Bond covalently with each other in family

c. Polonium (Po) / Tellurium (Te) are semi-metals

Halogens

a. Group 17 on periodic table; 7 valence electrons

b. React with alkaline to form salts

1. Salt is when – ion is given by acid and + ion is given by base

Noble Gases

a. Group 18 in periodic table; outer shell has 8 valence or is full

b. Inert

c. Found in free state in nature

d. Previously thought no compounds could be made from this family; 1962 Xenontrioxide (XeO₃) and Kryptondifluoride (KrF₂)

Lanthanide Series

a. Soft, malleable metals

b. High luster

c. Generally used to make alloy

Actinide Series

a. All are radioactive

E. Metals and Non-metals

Metals

a. good conductors of heat and electricity, malleable, ductile, lustrous, solid with exception of Mercury (Hg)

Non Metals

a. poor conductors of heat and electricity, dull appearance, not malleable, found as gas, liquids and crystalline solids like Iodine (I)

Semi-Metals and Metalloids

a. Stair step case

b. Have properties of both

F. Atomic Diagram

The space in which electrons orbit around the nucleus is the electron cloud
In the electron cloud, there are up to seven energy levels
Two ways to designate energy levels: 1, 2, 3, 4, 5, 6, and 7; or K, L, M, N, O, P, and Q
2n^2

a. 2e-, 8 e-, 18 e-, 32 e-, 50 e-, 72 e-, and 98 e-

With each energy level you will find 1 or more sublevels

a. These are designated by s-2 e- p-6 e- d-10 e- f-14 e-

Home

Notes- Compounds

I. Compounds

a. Two or more elements that are chemically combined

b. All physical and chemical properties are changed

II. Covalent

a. Formed between 2 nonmetals

b. Share electrons

c. Form molecules

III. Ionic

a. Ionic compounds must have at least 1 nonmetal and 1 metal.

b. Metals generally have 3 e- or less and lose electrons.

c. Nonmetals gain electrons.

d. Oxidation Number- the number of electrons that an atom will lose or gain

i. metals- positive- means they lose electrons

ii. nonmetals- negative- means they gain electrons

Home

Notes-Chemical Bonds

A. Chemical Formula

Introduction

a. (6.23x10²¹ molecules in water) Atoms in water are bonded (chemically joined together) to make identical water molecules

b. Regardless of state bonding ratio remains the same

c. Salt is ionic compound not molecule. Molecules are identical in salt crystal

d. Subscript one is assumed

e. Chemical formulas show:

1. Symbols of the elements in the compound

2. Number of atoms of each element (subscript)

B. Bonding-chemically joined

Covalent- two non-metals, share electrons until outer ring has 8 electrons; make molecules

a. Diatomic elements- two identical atoms bonded

b. Double bond means atoms are sharing two electrons each

c. H 1e ^–

H₂

HxxH

H-H

d. Sharing 1 electron each

Ionic- metals and non-metals. Metals lose electron to become positive ion and non-metal gains electron to become negative ion and non-metal gains to became a negative ion

1. MgS-ionic

2. SO₂-covalent

3. CCl₄-covalent

4. LiCl-ionic

5. KBr-ionic

6. NO₂-covalent

7. I₂-covalent

8. CH₄-covalent

C. Oxidation No.’s

Oxidation number is the charge that an atom acquires by gaining or losing electron in ionic bonding.
Metals will have positive. Non-metals will generally have a negative charge.
Oxidation numbers add up to 0 atoms.
Chlorine Magnesium 1 2 3

Oval: p+ 17
no 18 p+ 17

n^o18

e- 17

Cl^- 8e-

2e- 8e- ~~7e-~~

p+ 12

Oval: p+ 12
no 12 n^o 12

e^- 12

Mg²⁺ 2e- 8e- ~~2e-~~

8e-

Cl^- 2e- 8e- ~~7e-~~

D. Writing Formulas

The criss-cross rule (for ionic compounds only)

a. Mg²⁺ Cl^-

Chlorine Magnesium

Cl^- Mg²⁺

Mg²⁺ Cl^-

Mg Cl₂

MgCl₂

b. Oxidation for one element becomes subscript for other. (Only works for when oxidation numbers are different and cannot be reduced)

c. S^2- Ca²⁺

d. If you reduce Ca₂S₂à Ca_2/2S_2/2à Ca~~_2/2~~S~~_2/2~~àCaS

E. Naming Compounds

Binary Ionic- Stok naming system

a. Stok system: (for binary ionics)

b. Name the metal first

1. When metal has more than 1 oxidation number, you change the oxidation number to a Roman numeral, placed inside parenthesis, after the name of the metal

c. Name the nonmetal and add suffix “-ide”

d. Example: Sr₂₊S_2-

Strontium Sulfide

e. Example: CrCl₃

Cr_{2+, 3+, 6+}Cl_1-

Chromium (III) Chloride

Binary Molecular

a. Covalent bonding between two non-metals

b. You do not reduce subscript

c. Old System

1. 10 Prefixes- tell number of atoms

1. Mono

2. Di

3. Tri

4. Tetra

5. Penta

6. Hex

7. Hepta

8. Octo

9. Nona

10. Deca

2. Use only when both are non-metals

3. To the first non-metal never use “Mono-“

4. To the second non metal use prefix and add “-ide”

d. Example:

-CI₄

Carbon Tetraiodide

-SiO₂

Silicon Dioxide

-N₂O

Dinitrogen monoxide

-P₂O₅>

Diphosphorus Pentoxide

-As₂S₃

Diarsenic Trisulfide

Chemical Symbol	Type of Compound	Name
HgCl₂	Binary Ionic	Mercury (II) Chloride
Al₂O₃	Binary Ionic	Aluminum Oxide
N₂O₃	Binary Molecular	Dinitrogen Trioxide
MnCl₃	Binary Ionic	Manganese (III) Chloride
CBr₄	Binary Molecular	Carbon Tetrabromide
SnI₄	Binary Ionic	Tin (IV) Iodide
ClO₂	Binary Molecular	Chloride Dioxide
CuS	Binary Ionic	Copper (II) Sulfide

Ternary Ionic or Pseudobinary Ionic (contain polyatomic ions)

a. Polyatomic Ions

1. Groups of atoms that behave as a single atom. They come together covalently, but they are still not stable.

2. Polyatomic ions have a name, formula, and oxidation number (charge)

b. Example:

1. Na⁺ HCO₃^-

(sodium) (Bicarbonate)

NaHCO₃ (any name on common radical does not change)

2. Ca²⁺ CO₃^2-

CaCO₃

Calcium Carbonate

3. Fe^{2+, <3+>}PO₄^3-

FePO₄

Iron (III) Phosphate

c. Whenever you need more than one of the polyatomic ions, you place the formula of the polyatomic ion in paraenthesis.

1. Ca²⁺ClO₃^-

Ca(ClO₃)₂

Calcium Chlorate

2. NH₄⁺ S^2-

(NH₄)₂S

Ammonium Sulfide

Home

Notes- Weights-n-Percents

A. Weights

By using the atomic weight assigned to elements you can determine the formula or molecular weight of a compound

a. Formula weight (for ionic)

b. Molecular weight (for covalent)

c. The weight is determined by multiplying mass number times subscript for that element of the compound, then total the weight

d. Anhydrates (no water)

1. Strontium Hydroxide

Sr²⁺ OH^-

Sr(OH)₂

# of atoms mass no.

Sr 1 x 88 = 88

O 2 x 16 = 32

H 2 x 1 = 2

F.W 122

2. Sodium Bicarbonate

Na⁺ HCO₃^-

NaHCO₃

# of atoms mass no.

Na 1 x 23 = 23

H 1 x 1 =1

O 3 x 16 = 48

Cb 1 x 12 = 12

F.W 84

3. Arsenic Trichloride

AsCl₃

# of atoms mass no.

As 1 x 75 = 75

Cl 3 x 35 = 105

M.W 180

e. Hydrates- having water

1. Barium Chloride with two water molecules; Barium Chloride Dihydrate

BaCl₂ 2H₂O

Ba- 1 x 137 =137

Cl- 2 x 35 =70

H- 4 x 1 =4

O- 2 x 16 =32

F.W. 243

B. Percent Composition

To determine percent by mass of each element in the compound, you divide the total weight of the element by the total weight of the compound and multiply by 100. Percent must add up to 100.

a. Magnesium Phosphate

Mg ²⁺ PO₄^3-

Mg₃(PO₄)₂

Mg-3 x 24 =72 ⁷²/₂₆₂ x 100 =27%

P-2 x 31 =62 ⁶²/₂₆₂ x 100 =24%

O-8 x 16 =128 ¹²⁸/₂₆₂ x 100 =49%

FW 262 100%

b. Mercury (II) Nitrite

Hg ²⁺ NO₂^1-

Hg(NO₂)₂

Hg-1 x 201 =201 ²⁰¹/₂₉₃ x 100 =69% =68.6%

N-2 x 14 =28 ²⁸/₂₉₃ x 100 =10% =9.5%

O-4 x 16 =84 ⁶⁴/₂₉₃ x 100 =22% =21.8%

FW 293 101% 100%

c. Magnesium Nitrate Hexahydrate

Mg ²⁺ NO₃^1- - 6H₂O

Mg(NO₃)₂- 6H₂O

Mg-1 x 24 =24 ²⁴/₂₅₆ x 100 =9% =9.4%

N-2 x 14 =28 ²⁸/₂₅₆ x 100 =11% =10.9%

O-6 x 16 =96 ⁹⁶/₂₅₆ x 100 =38% =37.5%

H-12 x 1 =12 ¹²/₂₅₆ x 100 =5% =4.7%

O-6 x 16 =96 ⁹⁶/₂₅₆ x 100 =38% =37.5%

FW 256 101% 100%

To find the percent of Magnesium Nitrate in this hydrate

Mg-1 x 24 =24

N-2 x 14 =28

O-6 x 16 =96

148 ¹⁴⁸/₂₅₆x100 = 57.8% or 58%

H-12 x 1 =12

O-6 x 16 =96

108 ¹⁰⁸/₂₅₆x100=42%

C. Formula by Percent

Given that a compound is composed of a certain percentage each element you can find the formula of a compound
Example:

Ni 46% Cl 54%

Ni ⁴⁶/_{mass no. 59}Cl ⁵⁴/_{mass
no. 35}

Ni .780 Cl 1.54

Ni ^.780/_.780= 1 atom Cl ^1.54/_.780= 1.97 atom

Smallest number

NiCl₂; Nickel (II) Chloride

Example:

N 64% O 36%

N ⁶⁴/_{mass no. 14}O ³⁶/_{mass
no. 16}

N 4.57 O 2.25

N ^4.57/_2.25= 2 atom O ^2.25/_2.25= 1 atom

N₂O; Dinitrogen monoxide

Example:

Cu 52% C 10% O 38%

Cu ⁵²/_{mass no. 64}C ¹⁰/_{mass
no. 12}O ³⁸/_{mass no. 16}

Cu .813 C .833 O 2.38

Cu ^.813/_.813= 1 atom C ^.833/_.813= 1 atom O ^2.38/_.813= 3 atom

CuCO₃; Copper (II) Carbonate

Home

Notes- Chemical Equations

1. Law of Conservation of Mass (Matter)- Matter cannot be created nor destroyed, matter is conserved.

2. Word Equations

How would you read 2Mg+O₂à 2MgO

1. Magnesium and Oxygen yield Magnesium Oxide

2. + means “and”, and à means “produces or yields”

How would you read this equation? N₂+3H₂à2NH₃

1. Nitrogen and Hydrogen produce Ammonia

2. Common Acids and Bases:

a. HCl- Hydrochloric Acid

b. H₂SO₄- Sulfuric Acid

c. HNO₃- Nitric Acid

d. NH₃- Ammonia (not to be confused the Ammonium NH₄)

3. Chemical Equations

a. A chemical equation contains the formula of each chemical in the reaction

b. Coefficients are always written before the compound

4. Balancing Equations

a. Never change the subscript

b. Coefficients are in the front of element or whole compound

c. Always use lowest possible number of atoms

d. 2N₂+6H₂à4NH₃reduces to N₂+3H₂à2NH₃

e. No coefficient indicates one molecule

f. Chemical Equation: Cu+2AgNO₃à2Ag+Cu(NO₃)₂; Word Equation: Copper and Silver Nitrate yields Silver and Copper (II) Nitrate

g. Physical States is shown in paranthesis.

G	Gas
S	Solid
L	Liquid
Aq	Aqueous

h. O₂(g) can be called Oxygen gas or Gaseous oxygen

i. NaCl(s) is called solid sodium chloride

j. H₂O(l) can be called water, liquid dihydrogen monoxide, or even liquid water

k. HCl(aq) is called Aqueous Hydrochloric Acid

l. CaCl₂ ¯ - the “¯” indicates formation of precipatate, which is Calcium Chloride

m. H₂ - the “ ” indicates Hydrogen gas is released

n. à w/triangle over indicates “when heated prodcues”

4. Reactions

a. Chemical Equation: Zn(s) + 2HCl (aq) à ZnCl₂(aq) + H₂; Word Equation: Solid Zinc and aqueous hydrochloric acid yield aqueous Zinc Chloride and Hydrogen gas is released.

Zn 1 à 1

H 1x2 à2x1

Cl 2x1 à 2

5. Example Problems:

1) Write a balanced chemical formula for the following: Solid Aluminum and Oxygen gas when heated produces Solid Aluminum Oxide

4Al(s)+3O₂(g) àw/triangle 2Al₂O₃(s)

Al 1x4 à 2x2

O 2x3 à 3x2

2) Write a word equation and correctly balanced chemical formula equation for the following: CuSO₄(aq)+Fe(s)àFe₂(SO₄)₃(aq)+Cu(s)

3CuSO₄(aq)+2Fe(s)àFe₂(SO₄)₃(aq)+3Cu(s)

Aqueous Copper (II) Sulfate and Solid Iron yield Aqueous Iron (III) Sulfate and Solid Copper (Note: if polyatomic ion stays together treat as atom)

Cu 1x3 à 1x3

SO₄ 1x3 à 3x1

Fe 1x2 à 2x1

3) Write a balanced chemical formula for the following: Sodium Sulfate and Iron (III) Nitrate produce Sodium Nitrate and Iron (III) Sulfate

3Na₂SO₄+2Fe(NO₃)₃à6NaNO₃+Fe₂(SO₄)₃

Na 2x3 à 6x1

SO₄ 1x3à3x1

Fe 1x2à2x1

NO₃ 2x3à6x1

4) Write a balanced chemical formula for the following: Aqueous Calcium Carbonate and Aqueous Hydrochloric Acid produces Solid Calcium Chloride and Liquid Water and Carbon Dioxide gas is released.

CaCO₃(aq)+2HCl(aq)àCaCl₂(s)+H₂O(l)+CO₂

Ca 1à1

C 1à1

O 3 à 1+2

H 1x2à2x1

Cl 2x1 à 2

5) Write a balanced chemical formula for the following: Calcium Hydroxide and Ammonium Sulfate produces Calcium Sulfate and Water and Ammonia.

Ca(OH)₂+(NH₄)₂SO₄àCaSO₄+2H₂O+2NH₃

Ca 1à1

O 2+4à4x1+1x2

H 2+8à2x2+3x2

N 2à1x2

SO₄ 1à1

6) Write a balanced chemical formula for the following: Potassium Dichromate and Hydrochloric Acid and Tin (II) Chloride produces Tin (IV) Chloride and Potassium Chloride and Chromium (III) Chloride and Water.

7) Check to see if the following equations are correctly written, then balance. Mg₂SO₄+Na(PO₄)₃àNa₂PO₄+MgPO₄

8) Check to see if the following equations are correctly written, then balance and write the word equation. Ga+3H₂SO₄àGa₂(SO₄)+H₂

9) Write a balanced chemical formula for the following: Cesium Permanganate and Sulfuric Acid and Calcium Oxalate produces Calcium Sulfate and Cesium Bi-Sulfate and Manganese (II) Sulfate and Carbon Dioxide.

Home

Notes- Reactions

I. Synthesis

a. 2 or more elements or compound combining to form 1 compound

b. end up with 1 compound

c. 2Mg +O₂ à 2MgO

II. Decomposition

a. 1 reactant and 2 or more products

b. 2LiCl à 2Li +Cl₂

III. Displacement

a. Single

i. A+BCà B+AC

ii. A is more active than B

iii. Metal + Ionic Bond (B is metal C is nonmetal)à Metal (from Ionic Bond) +Compound (FreeMetal and Ionic Nonmetal)

iv. A (NM) +B(M)C(NM)àC(NM)+B(M)A(NM)

v. Zn+2HClà H₂+ZnCl₂

b. Double

i. AB+CDàAD+CB

ii. Compound + Compound à New compound + new compound

iii. BaCl₂ + Na₂SO₄à BaSO₄+ 2NaCl

IV. Combustion

a. Fuel combines with O₂ releasing large amounts of energy in the form of heat and light

b. Complete

i. Plenty of air present

ii. High temperature

iii. Carbon will be oxidized at its highest Oxidation number (+4)

iv. Whenever a fuel is burned you will always produce CO₂+ H₂O

v. 2C₈H₁₈+25O₂à 16CO₂+18H₂O

c. Incomplete

i. Limited amount of air

ii. Lower temperature

iii. Carbon oxidized at its lowest state (+2)

iv. Always produces CO + H₂O

v. 2C₈H₁₈+17O₂à 16 CO+18 H₂O

Home

Notes- Acid and Base Chemistry

Intro to Acids-n-Bases

How Do Bases react with Water?

a. All acid and base are ionic

b. Ionic compounds disassociate (ionize) in water

c. Base dissolves and one is always OH-

d. Mg(OH)₂àMg 2+ + OH-; Mg(OH)₂disassociates in water

How Do Acids react with Water?

a. Acid disassociates, producing H+ (hydrogen ions)

b. HClàH+ + Cl -; HCl disassociates, producing hydrogen and chloride ions.

Rules for Naming

Acids- start with H and contain non-metal

a. Binary

1. add prefix “hydro” to acids with Group VII elements

b. Ternary- polyatomic ion and H+

1. Rule 1- add name of ion plus “ic” for acids with polyatomic ion with poly atomic ion ending –ate

2. Rule 2- add name of polyatomic ion plus “ous” plus acid with polyatomic ion ending -ite

Bases- contain OH and usually contain a metal

a. Naming is same as naming ionic compounds

b. NH₃-Ammonia and NaCO₃ are bases without hydroxide

c. Examples for naming: Pb(OH)₂ – Lead (II) Hydroxide; Sr(OH)₂ – Strontium Hydroxide

Properties

Acids- Produce H+ when disassociates

a. Corrosive- cause a destructive chemical change in metals

b. React with more active metals to produce hydrogen gas

c. Edible acids are sour, some maybe poisonous

d. Acidic solution conducts electricity (electrolyte)

Bases- Produce OH- when disassociates

a. Bases are corrosive

b. Bases dissolves fat and oils

c. Edible bases are bitter, and like acids some are poisonous

d. Bases feel slippery (the kind that are safe to touch)

e. Bases conduct electricity (electrolyte)

f. An example of a common base is soap

Strength v. Concentration

Acids/Bases

a. Acids

1. Strong Acids

a. HCl is a strong acid composed of H and Cl molecules

b. A strong acid is one that completely disassociates in water

c. List of strong acids: Hydrochloric Acid- HCl, Sulfuric Acid- H₂SO₄, Nitric Acid- HNO₃,

2. Weak Acids

a. Acetic Acid is weak because it does not completely disassociates in water

b. List of weak acids: Carbonic Acid- H₂CO₃, Acetic Acid- HC₂H₃O₂, Citric Acid- H₃C₆H₅O₇

b. Bases

1. Strong Bases

a. A strong base is on that completely disassociates in water

b. List of Strong base: Potassium Hydroxide- KOH, Calcium Hydroxide Ca(OH)₂, and Sodium Hydroxide NaOH

2. Weak Bases

a. Weak base is one that does not completely disassociate

b. List of Weak base: Aluminum Hydroxide Al(OH)₃, Ammonia NH₃, and Magnesium Hydroxide Mg(OH)₂

2. Concentration of Acids and Bases

a. Concentration (has nothing to do with base)-refers to number of solute molecules per volume of solvent

b. For acid & base solution number of molecules per volume of water

pH Scale

Measures amount of H+ ions
Acid range 1-6.9, Neutal 7, and Base range 7.1-14
Hydrogen ion count increases closer to zero
Hydroxide ion count increase closer to 14
Neutral is a balance of H+ ions and OH- ions
Indicators- dye that shows one color in the presence of hydrogen, and another color in the presence of hydroxide
Red paper turns blue in base; Blue paper turns red in acid
Phenolphthalein-turns colors in presence of acid or base

Neutralization Reaction

Acid Base Showdown

a. Neutralization reaction occurs with an acid and base

b. 2Na⁺+H 2 SO42^-àNa₂SO₄

c. Neutralization reaction produces salt and water

d. A salt is formed when negative ion from an acid and a positive ion from the base combine

e. 2HNO₃+Ba(OH)₂àBa(NO₃)₂+2H₂O

f. NaHCO₃+HClàNaCl+H₂O+CO₂

Home

Physics

Notes- Force-n-Motion

(Physics Definition)- The science that deals with matter, energy, and their interactions

1. Matter- anything that has space and has mass

2. Energy- is the ability to do work

A. Force

Effect of Force on Motion

a. Force- action that can change the motion of an object (can force something to move, stop, speed up, slow down, change direction)

b. Forces are changing direction and speed (causing acceleration)

c. A force applied briefly is all that is needed to make an object move in a straight line (motion continues after force is removed)

d. A force applied constantly is needed to keep an object moving in a circle

B. Centripetal Force

Centripetal force is a force that pulls objects inward
Inertia- to resist a change in motion
Centripetal force makes motion circular

Earth Revolves around the Sun at	60,000 mph
Earth Rotates around its axis at	1,000 mph
Escape Velocity of Earth Gravitation is	25,000 mph

C. Gravitational Force-n-weight

Acceleration of gravity (g) = 9.8 meters/second/second or meters/sec²
This is if air resistance does not interfere with motion, gravity causes all objects to fall with the same acceleration regardless of mass.
Terminal Velocity- point where the object will stop accelerating because the upward force air resistance will equal to the downward force of gravity. On earth an average terminal velocity is 120 mph
More force is needed to hold up a heavier object than to hold up a lighter object
Gravity: measure of force of attraction between two objects
More mass greater gravitational force
Weight

a. Weight: Force of gravity on an object

b. Weight does not equal mass

c. Mass is a property of matter, but not weight

d. No gravity = No weight

e. Moon’s gravitational force is 1/6 of Earth’s, Mars is 1/3 gravitational force of Earth, Jupiter is 2.5 times the gravitational force of Earth

D. Momentum

Momentum: strength of tendency to keep moving
Same speed; more mass = greater momentum
Same mass; greater speed = greater momentum
p = m x v

p: momentum (kg x m/s)

m: mass (kg)

v: velocity (speed in a certain direction) (in m/s/s)

E. Energy- ability to do work

There are two types of mechanical energies:

a. Kinetic - energy of motion; depends on 2 things: mass and velocity

Remember: an increase in either mass or velocity will mean an increase in KE

b. Potential - Energy of position; object not moving- stored energy

Ex. A stretched rubber band has the ability to fly across the room

1. gravitational potential energy- it is potential energy that is dependent on height.

2. Kinetic - Potential energy conversions

a. Changes in the forms of energy are called energy conversion

b. Ex. A continuous conversion between KE and PE takes place in a pendulum

F. Newton’s Laws

Laws

a. Scientific Law - an accepted statement based on a large amount of evidence. A law can be disproved by new evidence, therefore subject to change.

1^st Law (Law of Inertia) - An object at rest remains at rest and an object in motion continues in motion unless acted upon by an unbalanced force.
2^nd Law (Law of Acceleration) - Acceleration of a body is directly proportional to the force exerted on the body and inversely proportional to the mass of the body and in the same direction of the force being applied.

a. Force = Mass x Acceleration; Force is measured in N

b. Newtons= kg x m/sec²

c. 1 lbs. = 4.45 N

d. Example Problems:

1. A racecar driver wishes to accelerate at 15 m/s² in a car that weights 750 kg. How much force must be applied by the wheels to the ground?

1) F=M x A

2) F=750kg x 15 m/s²

3) 11,250kg x m/s²

4) 11,250 N à(3 sig fig)à11,300 N

2. A sixty kilogram person riding on a fifteen kilogram sled is pushed with a force of 300N. Find the person’s acceleration.

1) F=M x A

2) 300 kg x m/sec² = 60 kg + 15 kg

3) (300 kg x m/sec²)/ 75 kg = ~~75 kg/ 75 kg~~

4) 4.00 m/sec² = Acceleration

3^rd Law (Law of Interaction) - For every force or action there is an equal and opposite force or reaction.
Newton’s Law of Universal Gravitation

a. A force of attraction exists between every 2 objects in the universe

b. The greater the mass, the greater the gravitational force.

c. The farther apart the objects, the smaller the attraction.

Home

Notes- Work, Power, & Machines

I. Work

a. A force through a distance the object moves must be in the same direction the force is applied

b. Work=Force x Distance

c. Force- Unit is Newton, or kg m/sec²

d. Work- Unit is Newton-meter or Joule

II. Power

a. Power=^{Force x Distance}/_Time

b. Power= Work/Time

c. Unit is Joule/sec, or Watt

d. 745.56 Watts= 1 Hp

III. Machines

a. Make Work easier by

i. Change the amount of force applied

ii. Changing the distance

iii. Changing the direction of the force

b. Work Input- force put into machine

i. E.F. x E.D.

ii. Force x Length

c. Work Output- what machine does

i. R.F. x R.D.

ii. W(weight) x H (height) (Inclined Plane)

d. Efficiency- work output compared to work input (always a percent)

i. W.O./W.I. x 100

ii. If no friction, WO=WI

e. Mechanical Advantage- Number of times the machine multiplies the force

i. AMA= Actual Mechanical Advantage (including friction)

ii. IMA= Ideal Mechanical Advantage (No Friction)

IV. Simple Machines

a. Inclined Planes

i. Flat slanted surface

ii. AMA= R.F./E.F.

iii. IMA= L/H

b. Wedge

i. 2 inclined planes that move

ii. The longer and thinner the wedge, the less the effort applied

iii. To increase Mechanical Advantage of a wedge, you sharpen it

c. Screw

i. An inclined plane wrapped around a central bar

ii. To find the IMA, you count the number of threads per inch

iii. More threads per inch increase the mechanical advantage

d. Lever

i. Rigid bar that is free to move about a fixed point

ii. Fulcrum- pivot point

iii. Resistance- weight

iv. Effort force- what you push with

v. Arms

1. Resistance Arm- distance from weight to fulcrum

2. Effort arm- distance from effort force to fulcrum

3. When E.A. is longer, the machine will increase the force

a. Mechanical Advantage is greater than 1.

4. When R.A. is longer, the machine will increase the distance

a. Mechanical advantage is less than 1.

vi. AMA= R.F./E.F.

vii. IMA= E.D./R.D.

viii. 3 classes

1. 1^st Class Lever

a. R.F. ____________F______________E.F

b. Always changes the direction of the force

2. 2^nd Class lever

a. F_____________R.F.___________E.F

3. 3^rd Class Lever

a. F________E.F._________R.F.

b. R.D. is always longer that the E.D

c. Baseball bat

e. Pulley

i. Rope or belt wrapped around a wheel

ii. Can either change direction or multiply force

iii. AMA=R.F./E.F.

iv. IMA= Number of Supporting Ropes

f. Wheel and Axle

i. 2 circular objects of different sizes

ii. Wheel always large than axle

iii. Wheel goes greater distance then the axle

iv. IMA= Radius of Wheel/ Radius of Axle

Home

Notes- Heat Energy

KE and Temp.

The more energy molecules absorb, the greater their kinetic energy
Temperature- measure of the average kinetic energy of molecules
Scales for Temperature

a. Centigrade Scale- 0°C water freezes, 100°C water boils

b. Fahrenheit Scale- 32°F water freezes, 212°F water boils

c. Kelvin - 0K is absolute zero (theoretical temperature which molecules stop moving) freezing point of water is 273K and boiling point is 373K

1. Absolute Zero= -273°C, -460°F, 0°K

2. Closest ever at University of Colorado - Rubidium (Rb) less than ¹/_{70,000,000,000} K

Conversions Formula

a. F = 1.8x°C+32

b. C=(F°-32)/1.8

c. K=C°+273

Heat Formula

a. Heat - is a form of energy caused by internal movement of molecules; deals with two things, number of molecules and velocity

b. Heat Content- total heat content in a substance, measured by calories or joules

c. Heat Content = Mass (g) x ∆T, °C x sp. H

Energy Transfer

Conduction - Molecules transfer energy by physical contact; in solid state of matter
Convection - Molecules carry energy from one place to another; in gas and liquid states
Radiation - energy transferred without direct contact
Conduction = touch, Convection = rise, Radiation = everywhere

Heat Flow

1. Heat flows from higher temperature to lower temperature

Specific Heat

Everything has specific heat
Amount of heat needed to change temperature of a substance depends on its mass
Specific Heat is:

· How difficult something is to heat or cool?

· Measure of heat needed to raise the temperature of a substance by a certain amount

· Units for sp. H: ^cal/_g°_C and ^J/_g°_C

Higher the number, harder to heat, lower the number easier to heat
Specific Heat of water = 1 ^cal/_g°_C

Specific Heat of iron = .1 ^cal/_g°_C

Specific Heat of silver = .05 ^cal/_g°_C

Long heating time, Long cooling time
Expansion - molecules absorb energy, greater kinetic energy, move faster, take up more space
Contraction- molecules lose energy, less kinetic energy, move less, take up less space
3 substances that do expand when cooled are: Water (which begins expanding at 4°C), Bismuth (Bi), and Antimony (Sb)

State Changes

All particles of matter possess kinetic energy. As the temperature of a solid substance is raised the kinetic energy of the particles increases up to a certain point and then stops. At this point, the solid is melting. A change from one form to another, for example a liquid to a solid, or a solid to a gas, is known as a change of phase or a change of state. At this temperature, the particles have so much kinetic energy that they are just about to break their links. If more energy is added some of the particles do break their links and turn into a liquid. So all additional energy goes into link breaking, not kinetic energy. Since temperature measures kinetic energy, the temperature stays constant until all of the substance is melted. Then the particles can gain kinetic energy again, so the temperature starts to rise. This happens again as the temperature is raised to the boiling point.
The particle of a solid at a temperature below the melting point have less kinetic energy then particles of the same substance above the melting point. Likewise, particles of a liquid below the boiling point have less kinetic energy than particles of the same substance above the boiling point. The state of matter of a substance depends on its temperature, and higher temperature depend on higher kinetic energies.

Coefficient of Linear Expansion

Linear deals with solids
Def - the change in a unit of solid when its temperature is changed by 1°C. Remember: All solids expand and contract differently from each other.

Coefficient of Volume Expansion

Deals with liquids
The change in unit of a volume of liquid when its temperature is changed by 1°C. Remember: All liquids have different coefficient number
Gases: Charles’ Law - a gas expands by ¹/₂₇₃ of its volume at 0°C for each °C rise in temperature

Specific Heat Formula for Metals

a. SpH_m=M_w x SpH_w x DT_w / M_m x DT_m

Home

Notes-Waves & Sound

Waves

Wave – A disturbance that moves from one place to another.

*Crest – moving high point

*Trough – moving low point

*Amplitude – measure of wave intensity; distance from midpoint to crest.

*Wavelength – Distance from crest to crest; symbol: (Lambda)

Frequency, Speed and Wavelength

Frequency – the number of crests passing a point in a given time. (Time is always in seconds)

*Unit for frequency: Hertz (Hz)

*Hz = /sec (“per second”)

*Hz = crests/second (“crests per second”)

1: A wave has 70 crests in 35 seconds; find the frequency.

70 (crests) / 35 seconds = 2.0 /sec. or 2.0 Hz.

Speed of Waves:

Speed = Distance / Time

1: A wave moves 10 meters in 30 seconds. Find its speed.

S= D/T

S = 10 m / 30 sec.

Speed = .33 m/sec.

Higher frequency waves will always have the shorter wavelengths.

Lower frequency waves will always have the longer wavelengths.

When speeds are equal, wavelength and frequency are inversely proportional.

When wavelengths are equal, speed and frequency are directly proportional. Higher speed = Higher frequency; Lower speed = Lower Frequency.

Frequency = Speed / Wavelength

1: A wave has a speed of 30 m/sec and a wavelength of 2 m; calculate

the frequency.

F = S/D F= ^30m/sec/_2m= 15/sec = 15 Hz.

2: A sound wave of 6,800 Hz has a wavelength of .05 m, find the speed.

S = fD S =6,800Hz * .05m

S = 6,800 /sec * .05 m= 340 m/sec.

3: Find the wavelength of a wave with a speed of 343 m/sec and a frequency of 131 Hz.

D= S/F D = ^343m/sec/_131Hz

D = ^343m/sec/_131/sec

D= 2.6 m

Types of Waves

*Transverse Wave – motion of particles is perpendicular to the path of the wave; does not need a vacuum; light is a transverse wave; always travels in a straight line.

*Compression Wave (Longitudinal Wave) – motion of particles is parallel to the path of the wave; cannot travel in a vacuum; needs a medium; temporarily changes the density of the medium (displacement). The Crest (compression) has the highest density The Trough (rarefaction) has the lowest density; sound is a compression wave; sound waves move in all directions.

Wave Interference

*Constructive Interference – increased amplitude when crests of different waves overlap. (Louder sound)

*Destructive Interference – decreased amplitude when crests of one wave overlap the troughs of another wave. (Softer sound or sound discontinued all together)

*Acoustics – the science of sound; acoustical tiles absorb sound.

*Dead Spot – a spot where no sound is heard.

Speed of Sound

Sound waves move faster through solids; the closer the particles, the faster the waves.

The speed of sound depends on 3 things:

*The density of the object or medium

*The elasticity of the medium.

*Temperature or the object or medium; the speed of sound is constant when temperature doesn’t change.

**Speed of sound in air: 331 m/sec.**

For every 1°C it gets warmer, sound will go .6 m/sec faster

For every 1°C it gets colder, sound will go .6 m/sec slower.

°C * .6m/sec + 331 m/sec.

Sound Phenomenon and Application

*Reflection: Sound waves bounce off barriers; a reflection of sound we can hear is an echo.

*SONAR: Sound Navigation And Range; gives off beeps; uses refection and constancy of the speed of sound to find ocean depths.

Sonar depends on reflection and constancy of speed.

Speed of sound in seawater: 1500 m/sec. Reflection time: .5 sec

*Ultrasound: produce sonograms; sonogram: high frequency sound produces greater detail. Ultrasound depends on reflection and constancy of speed.

*The Sound Barrier and Supersonic Flight:

A plane causes a wake of compressed air in its path and a sonic boom occurs when the wake of compressed air reaches the observer.

1^st person to travel above the speed of sound: Chuck Yeager in October 1947.

Mach 1: 741 mi/hr (speed of sound)

Mach 2: 1,482 mi/hr (2x the speed of sound)

Mach 3: 2,223 mi/hr (3x the speed of sound)

A sonic boom depends on the constancy of speed.

*Surgery with sound:

Ultrasound – shatters kidney stones; calcium deposits.

Surgery with sound depends on constructive interference & the constancy of speed.

Pitch and Intensity

*Pitch and frequency are directly proportional. The shorter the wavelength, the greater the frequency and the higher the pitch.

*Intensity: the greater the amplitude, the greater the intensity; intensity means volume. The higher the amplitude, the louder the sound; sound intensity is measured in decibels. (dB)

3 Ranges:

Sub-Audible: below 0 dB

Audible: 0 dB to 120 dB

Harmful: above 120 dB

--Long-term exposure of 90 decibels will cause permanent hearing damage.

*Hearing sounds:

Parts of the ear:

1. The outer ear – collects sound waves; sound waves travel to the eardrum; the eardrum vibrates; the eardrum separates the middle ear from the outer ear.

2. The middle ear – 3 small bones located in the middle ear; the hammer, anvil and stirrup; they reinforce the sound.

3. The inner ear – cochlea located here; filled with nerve cells; sound travels to the brain.

Sounds We Can And Cannot Hear

Infrasonic (Subsonic): below 20 Hz

Audible to humans: 20 Hz to 20,000 Hz

Ultrasonic: above 20,000 Hz (not audible)

A dog can hear from 15 Hz to 50,000 Hz.

Resonance

An object that is vibrating its natural frequency can cause a nearby object to start vibrating if the object has the same frequency. The ability of an object to vibrate by absorbing energy is called resonance.

*Tacoma Bridge: November 7, 1940: fell apart due to resonance because the frequency of the wind was the same as that of the bridge.

Home

Notes- Light Energy & EMS- Natalie

Light Energy

· Speed of light = 300 million miles per second, needs no medium to travel through

· Travels through light a million times faster

· 8.3 minutes for light to reach earth

· can travel through air, water and glass

· speed of light is constant within a given substance

· When light is in the same medium it keeps the same speed of light, if it goes into another medium it can speed up or slow down.

· D = S x T

· Moon = 390,000 Km from earth

· Light waves have speed, frequency and wavelength

· Light waves are = Transverse waves

· C = Celeritous (speed)

· C = Speed x Wavelength

· When speed is constant frequency and wavelength are inversely proportional (When frequency goes up wave length goes down )

· Frequency is ALWAYS constant

~~~~~~~~~~

· An atom absorbs energy when it gains an energy level, it becomes unbalanced and falls back down to its original energy level. This causes a release of a photon.

· Photon is a bundle of energy that determines the speed of light produced and that is why frequency is always constant.

· When speed of light decreases wavelength decreases.

· When speed of light increases wavelength increases.

· Constancy of speed and wavelength are directly proportional

Refraction

· It is the bending of light as it travels from one substance to another when entering at an angle.

· If the light rays hit another medium strait on the light ray does not refract it only slows down when entering another medium

· Incident Ray = The light ray that enters the substance at an angle

· Refracted Ray = The light ray exits the substance at an angle and speeds up or slows down depending on the substance

· The Normal is 90 degrees from the surface

· Angle of Incidence = Distance from the normal to the Incidence ray

· Angle of refraction = Distance from the normal to the Refracted ray

When light travels from a Less to a More Dense area

1. Ray speeds up (because light travels faster in more dense particles such as solids)

2. Refracted ray bends towards the normal

When light travels from a More to a Less Dense area

1. Ray slows down (because light travels slower in less dense particles such as air)

2. Refracted ray bends away from the normal

Lenses

· Index of refraction = Degree to which a substance bends light. (Higher the index number the more light is going to bend and slow down)

· Index # = (speed of light in a vacuum) / ( speed of light in that substance)

· Converge = Rays come together

· Diverge = Rays spread apart

· Focal Point = Point at which parallel rays meet after refracting

· Real Image = Image formed from the inverted light rays. The image is flipped and smaller. The image can be seen by the human eye and can be projected on a screen

Convex Lenses (Caves Out)

· Inside Focal Point Rays Diverge

· Outside Focal Point Rays Converge and produce a Real Image

· Formula = (do/di) = (so/si)

· Do = Object distance, Di = Image distance,

· So = Object size, Si = Image size

Concave Lenses (Caves In)

· Inside Focal Point Rays Diverge

· This Lens will always Diverge light and produce a virtual image

Virtual Image = An image that can NOT be projected onto a screen because the light from the object does not meet there. (It only appears to)

Nearsightedness

· Eyeball is too Long or too Convexed

· The Rays converge to make a clear image b4 the retina

· Also called Myopia

· Can be corrected by a concave lens which diverges the rays (causing the rays to move farther back and hit the retina)

Farsightedness

· Eyeball is too Short or Not Convexed Enough

· The Rays converge to make a clear image after the retina

· Also called Distopia

· Can be corrected by a convex lens which converges the rays (causing the rays to move closer together and hit the retina)

Mirrors

· A mirror is a smooth piece of glass with a sliver sheet in one side

· Reflection = Light reflects off everything, not just mirrors

· Reflection of light allows us to see them.

· Diffused Reflection = When a surface is Rough, the rays come in at one direction and bounce off in many directions.

· Regular Reflection = When a surface is Smooth, the rays come in at one direction and bounce off (angled) in another direction but together the same way.

· Luminous = An object that gives off its own light (sun).

· Illuminate = An object that reflects light (moon).

· Law of Reflection = Angle of Incidence is EQUAL to the Angle of Reflection (or the image will be distorted)

Concave Mirrors

· Inside Focal Point Rays Diverge

· Outside Focal Point Rays Converge and produce a Real Image

Convex Mirrors

· Inside Focal Point Rays Diverge

· This Lens will always Diverge light and produce a virtual image

Plane Mirrors

· Are Flat Every Day Mirrors that reflects the opposite of what is being mirrored (if u lift up your right arm in a plane mirror it looks as if you are lifting up your left in the mirror)

· Always Produce a Virtual Image

Visible light

· ROY-G-BIV

· Depending on the Frequency light changes color

· A Prism can be used to spread out visible light. (Triangular shaped piece of glass)

· When light passed through a prism the angle of refraction changes slightly with wavelength.

Example

- Blue light with the shortest wavelength has the highest frequency, is refracted the most.

- Therefore Red light with the highest wavelength has the shortest frequency is refracted the least.

(Refer to Objective 4: Light and Color Handout)

Higher Frequency the more light will refract!

Colors

· White light All the colors reflected

· Black light All the colors absorbed

· Clear Doesn’t absorb anything

Color is reflected to the eyes and absorbed by the Retina

Opaque = Light can NOT pass through and you can Not see through. All light is absorbed. (Black Shirt)

Transparent = Light CAN pass through and you Can see through. (Transparency)

Translucent = Light Can pass through but you can NOT see clearly (Frosted Glass)

Home

Notes- Electricity

1. Static Electricity

Refers to all of the property of charges that are not moving
Free electrons are everywhere.

2. Law of Electrical Charges

Unlike charges attract (+/-) and like charges repel
Atoms become positive when they lose an electron. Atoms become negative when they gain an electron.
Electroscope is used to detect electric charges

3. Static Electricity in the Atmosphere

Water droplets rub against each other (friction) and cause static electricity to accumulate
Positive ions go to the top and negative ions go to the bottom of cloud.
Induction- bringing a charged object near but not close enough to touch, causes negative ions on the earth to move away.
A leader emerges from the cloud and a streamer emerges from the ground.
Electrons from cloud become ionized and producing so much heat it causes atoms to glow. The air around the lightning bolt is heated, expands, and returns to its normal size. Cooler air fills this gap left by heated air and causes thundering noise.
Grounded- electrical contact with the ground (on most homes it is near electric meter box)
Electrons travel at 1/3 the speed of light

4. Current Electricity

a. Electrons moving steadily in a definite direction through a material.

5. Conductors and Insulators

Conductor- material through which electrons flow freely
Insulator- material that greatly resists the passage of electrons through them
Factors for a good conductor include length of wire, diameter of wire, type of metal etc.etc.etc.

6. Voltaic Cell

Alessandro Volta- produced the first steady flow of electric current by chemical means.
Also known as a wet cell
He took and electrolyte (a liquid that conducts electricity, in his case Sulfuric Acid) and placed to differently charged poles inside [Zinc(-) and Copper(+)]. He connected the poles to a light bulb. And the continuous movement of particles from one pole to another caused the light to turn on.

7. Electrodes

Name of the poles
Cathode (-)
Anode (+)

8. Dry Cell

Two or more are known as clpbatteries
Casing is usually made of Zinc and is the Cathode.
Carbon Rod which is made of graphite and functions as the anode.
Electrolyte allows electricity to keep passing through the cathode and anode.

9. Series

+ - + - + - + - + - + - + (You get the point).
They are used to increase voltage
Generally 1.5 volts each
Add the volts together to get the flow of electrons.
Current = Flow of electrons ( Voltage stays the same)

10. Electrical Circuits

Needs a source of electrons Such as in a dry cell, battery, or generators
Needs a wire to travel through
Needs a force to push the electrons through (Voltage) or Pressure
Any type of applianc
Switch (not necessary but is more convenient

Sources of Electrons

a. Examples = generators, power plans, wet and dry cells, and a car battery

11. Plugs

a. Standard 2 prong = you can plug them in different was and is hazardous because even if it is off and plugged in opposite it still had the ability to shock you.

b. Standard with Hot(Electrons enter and is the smaller prong) and a Neutral side

(This prong is larger and is the side that allows the electrons to leave the appliance). This type of plug is more useful because it is safer and is totally shock proof.

c. 3 Prong = 3 way dedicated grounding system. Is the Best plug because of grounded prong to absorb extra electrons and shock.

12. Loads

a. Any electrical appliance that you can plug into an electrical source.

b. Anything adding resistance to an outlet.

13. Switches

a. Convenient way to open and close a circuit.

b. Open = Off

c. Closed = On

14. Fuses and Circuit Breakers

a. Designed to break an overloaded circuit so that your appliances won’t burn down or overload.

15. Fuse box

a. Old Boxes have a have a meter with a wire that melts and needs to be replaced if there is an overload.

b. New Boxes

1. Some are Magnetic Trip, in this type of circuit breaker, the magnetic connection is pulled up by an electromagnet and causes the circuit to open and shut off.

a. These don’t need to be replaced, they only need to be switched back on after fixing the problem.

2. Some are Thermal Trip, in this type of circuit breaker, there is a piece of metal that connects the circuits and if it gets to hot it begins to bend and the connection is shut off.

16. Short Circuit

a. A sudden drop in resistance, which makes current go up.

b. When insulation in wrapped wires goes away, electrons always travel through the shortest path and cut through the spaces where there is no insulation and this causes a drop in resistance.

17. Ground Fault Interrupters

a. These work on the principle of magnetism.

b. 31 thousandths of a second in response to .004 to.006 amps.

c. Outlets have three prongs and also a test and a reset button.

18. Factory Affecting Electric Current

a. Ampere is the measure of amps, 10 in 1 sec

Amp = 6.3 x 10^18 of an electron discovered by Andre Ampere

19. Electro Motive force

Difference of potential
Voltage = force pushing electrons
Everyday home volts = 120v – 125v
Large Appliances = (Electric Dryers) 240v – 250v

20. Resistance to flow

a. Anything that’s opposing the flow of electrons

b. Units of Ohms

c. Ohm meter

d. IF resistance gets too high it wont work because of insulators

21. Ohms Law

Current and voltage are directly proportional.
Current and resistance are inversely proportional

22. Circuits

1. Series = where the electron flow is along a single path and any break in this pattern will stop the flow of electrons.

Characteristics of a Series

a. Sum of the voltages across each load is equal to the voltages at the source

b. Current is the Same throughout the circuit

c. Total resistance is equal to the sum of all resistances in the circuit.

1. Parallel = Where the electrons have more than one path in which to travel through. And any break in 1 path will not stop the flow of electrons to the ofher paths

2. Characteristics of a Parallel

a. The Total Resistance will always be < than the smallest fingle resistance.

b. Voltage is the same throughout the circuit. It does not change 120v – 125v.

c. Total current flowing through the circuit = to the sum of all the current in the branches

23. Power and Energy

1. Alternating Current that enters homes Produced by generators.

2. US electrical frequency is 60Hz which changes direction 120 times/sec.

3. Each electrical appliance uses power at its own rate which is measured in watts

4. Formula for power = voltage x current(I) in amps

5. We have electrical meters in our homes that measures Electrical Energy.

6. Formula for Electrical Energy = power x time

7. EE = KwHr’s

Home

Notes- Magnetism & Electromagnetism

Lodestone – natural magnet.

Magnetism – property of a substance to attract another substance.

Theory of Magnetism – When more electrons spin in one direction the atom then takes on a north and south pole. When enough atoms align this causes the substance to take on a nort-south pole.

Domains – atoms aligned to north or south.

Ferromagnetic – something that is strongly attracted to a magnet ex. iron and cobalt.

Paramagnetic – slightly attracted to a magnet ex. sodium and platinum.

Diamagnetic – slightly repels from a magnet ex. mercury and bismuth

Nonmagnetic – allows magnetism to go through but it will not become magnetized ex. paper.

Types of Magnets

Magnetite – the material that makes up lodestones.

Alnico – consists of aluminum, nickel, cobalt, iron; a very powerful magnet.

Make a Magnet

There are three ways to make a magnet.

1) By contact – stroke a material the same way repeated times.

2) Induction – place a material near a magnet but they do not touch.

3) Electricity – Hooking up a solenoid to an electric current.

Hans Christian Orested discovered electromagnetism.

Electromagnet – temporary magnet.

Demagnetizing a Magnet

By heat – if placed in a flame and heated until it reaches the Curie point.

Curie point – every element that can be magnetized has a Curie point. It is a temperature where it loses all magnetic properties. Ex. iron – 770C cobalt – 1,113 C

By contact – stroked back and forth in order to disrupt the domains.

By hammering – by dropping or hammering a magnet which will cause the domains to become scattered thus demagnetizing it.

Law of magnetic poles – unlike forces attract; like forces will repel.

Flux lines – lines of fore in a magnet.

William Gilbert – person who said the earth behaved like a magnet.

Magnetic North Pole – located Hudson Bay, Canada.

Magnetic South Pole – located South of Australia.

Compasses go towards the Magnetic Poles.

Angle of Declination – angle between Magnetic South Pole, Magnetic North Pole and Geographic North Pole and the Geographic South Pole.

Electromagnets

Electromagnets differ from permanent magnets in three ways.

1) You have the ability to make the electromagnet stronger or weaker. You can have a thicker core, increase the current, more turns off wire, or the type of core.

2) You have the ability to turn the magnet off and on.

3) You are able to change the polarity by switch the direction of the current.

Ways We Use Magnets -Telephones, doorbells, motors, geenerators, and transistors.

Current, resistance, voltage in circuits

Andre Ampere – French scientist who named the ampere. (Amps)

6.3x10^18 – number of electrons per 1 Amp.

Volta – Italian physicist who named the volt.

Voltage makes electrons flow. It is measured with a voltmeter.

George Ohm – German physicist who named the ohm.

Resistance is measured in ohm. It is measured with an ohmmeter

Ohms Law

1) current and voltage are directly proportional to each other.

120 – 125v is common household voltage

2) current and resistance are inversely proportional

*** Formula***

Voltage (Volts) = Current (Amps) * Resistance (Ohms)

Series and Parallel circuits

The Series Circuit

The Series Circuit – where electrons flow through one path and any break will stop the flow of electrons.

Total voltage of the circuit will stop equal to the voltage of the load.

Current stays the same.

Total resistance will equal to the sum of the resistances in the circuit.

The Parallel Circuit

The Parallel Circuit – electrons travel in more than one pass; any break will not stop circuit.

Voltage stays the same. (120 -125v in a common household)

Total current flowing through the circuit is equal to the sum of the current in all of the branches.

Total resistance is always less than the smallest single resistance.

Powers and Energy

The current that enters your home is an alternating current.

Every appliance uses power in watts.

***Formula***

Power = Voltage * Current

Electrical Energy = Power (kilowatt) * Time (hours)

Electric meters measure electrical energy

Electric Motors

Electric Motors – change electrical energy into mechanical energy.

The Direct Current Motor

On a Direct Current motor the armature has two split rings called commutators. Do not need the commutators for a Alternating Current motor.

Electricity from a magnet

Michael Faraday – said you could get electricity from magnetism.

Electromagnetic induction – magnet or coil of wire must be moving to cut through the line of force. You create and induced current to increase induce current.

There are three ways to do this.

1) speed up motion (magnet or coiling wire)

2) get a stronger magnet

3) increase solenoid coiling

Generators – devices that produce electric current by electromagnetic induction. Converts mechanical energy into electrical energy.

Transformers – used to change the voltage produced at the power plants into our homes. Have a core and around each core is a coil of wire. When electricity comes in it’s through primary coil. When it goes out it is in secondary coil.

1) step up transformer – increases voltage

2) step down transformer – decreases voltage

Home

Hosted by www.Geocities.ws