Gateway Chemistry Review (Answer Key)
Structure and Properties of Matter
Matter

Anything that has mass and takes up space

All matter is made from three basic particles o Protons o Neutrons o Electrons

Protons, neutrons, and electrons make up atoms

Different types of atoms are called elements

Elements contain protons, neutrons, and electrons in differing numbers
Subatomic Particles

Nucleus
Contains protons and neutrons
Atomic mass is concentrated in the nucleus o Proton

Positively charged
 Found in the nucleus

Determines identity of the element

Mass = 1 amu o Neutron

Neutral

Found in Nucleus

Mass = 1 amu

Electron Cloud o Electron cloud surrounds nucleus o Contains particles which are negatively charged o Electrons are located at specific energy levels. o If the atom is neutral, the number of electrons equals the number of protons o Very small mass ( negligible )

1800 electrons equal the mass of one proton or neutron o Electrons in the outermost shell are called valance electrons.
Ions

An atom or group of atoms that has a positive or negative charge.

If an atom loses an electron, it becomes positive

If an atom gains an electron, it becomes negative .
Compounds

A substance containing atoms of more than one element
–
NaCl
– C
6
H
12
O
6
–
H
2
SO
4
–
C
13
H
18
O
2
(ibuprofen)
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Molecules

Two or more atoms bound so tightly that they behave as a single unit .

Linked by covalent bonds

Consist of atoms of the same element or different elements
Ionic Compound

Formed by the attraction of two ions that are oppositely charged.

Na
+
+ Cl

NaCl
Practice
Identify each of the following as an atom, ion, or molecule:

Ne atom

Cl
ion

Ca
2+ ion

CH
4 molecule

NO molecule

P
3ion
Density

Describes how closely packed atoms and molecules are in a given substance.
 The ratio of an object’s mass to its volume.

Volume of a cube = length x width x height

Density = mass/volume

Units: g/cm
3

Common Densities o Air: .001 g/cm 3 o Water (4
0
C): 1.00 g/cm
3 o Water/Ice (0
0
C): 0.92 g/cm
3 o Aluminum: 2.7 g/cm 3 o Gold: 19.3 g/cm
3
Practice:

CO


He
SO
2 molecule atom
4
2ion
1. Which object has a lower density, a brick or a block of Styrofoam? Styrofoam
2. Which object will float in water, a rock or a piece of ice? Why? Ice, it is less dense than water; a rock is
more dense than water.
3. What is the density of a substance that has a mass of 55g and a volume of 11cm
3
? 5g/cm 3
Mixtures and Solutions
Pure Substance

A type of matter in which all particles are of the same chemical composition o Au element o H
2
O compound o NaCl compound o C
6
H
12
O
6 compound o Ar element

Which of the previous examples is a compound? An element?

Why is salt water not a pure substance? Salt and water are two different substances mixed together in solution
Mixture

Two or more pure substances physically mixed together.

Cannot be represented by a single chemical formula. o Salt water
Which of the boxes below contain a substance pure
?
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o Sand and rocks o Air
Heterogeneous Mixture

A mixture where substances are not evenly distributed (non uniform) o oil and vinegar salad dressing o vegetable soup o sand and sugar o soil o granite
Homogeneous Mixture

A mixture where all components are evenly distributed (uniform).
 “same throughout” o salt water o gasoline o syrup o air
Solution

Formed when one substance is dissolved by another.

In order to be dissolved, a substance must be soluble.

A homogeneous mixture.

Particles are evenly distributed.

Parts cannot be separated by filtering .

Solvent— does the dissolving

Solute— dissolved by the solvent

Identify the solute and solvent in each of the following: o Salt water

Salt is the solute . Water is the solvent . o iced tea

Tea is the solute . Water is the solvent . o kool aid

Water is the solvent . Kool aid mix is the solute . o paint/paint thinner
 Paint is the solute . Paint thinner is the solvent . o nail polish/acetone (nail polish remover)

Nail polish is the solute . Acetone (nail polish remover) is the solvent .
Types of Solutions

Solid dissolved in a liquid. o Salt water

Gas dissolved in a liquid o Coca-Cola

Two solids o Metal alloys: brass = copper + zinc

Two gasses o Air: nitrogen (78% vol), oxygen (21% vol), argon (1% vol), carbon dioxide (0.03% vol).

In solutions of two solids or two gases, the solvent is the component present in largest quantity.
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Water
 The “ universal solvent
”

A solution in which water is the solvent is called an aqueous (aq) solution.

Does NOT dissolve everything . o Why is this a good thing?—think about the paint on your house.

Because water is polar, it dissolves other polar substances. o “ Like dissolves like ”

Water dissolves many other compounds.
Solubility

How much of a solute will dissolve in a given solvent.

How do you increase the solubility of a solid in a liquid? (hint: iced tea)
Heat the liquid.

How do you increase the solubility of a gas in a liquid? (hint: can of soda)
Cool the liquid.

Solid in a liquid o Increasing temperature will make a solid more soluble in a liquid. o Decreasing temperature will make a solid less soluble in a liquid. o Heat water before adding tea/sugar for iced tea.

Gas in a liquid o Increasing temperature will make a gas less soluble in liquid. o Decreasing temperature will make a gas more soluble in a liquid. o Increasing pressure will make a gas more soluble in a liquid. o Decreasing pressure will make a gas less soluble in a liquid.
Types of Solutions

Saturated o Holding the maximum solute at a given temperature. o “The perfect glass of sweet tea”; one more grain of sugar will sink to the bottom.

Unsaturated o Holding less than the maximum solute at a given temperature. o Lightly sweetened tea; additional sugar will dissolve

Supersaturated o Holding more than the maximum solute at a given temperature. o Overly sweet tea; sugar is forced into solution by heating and then cooling.
Questions

What term is used to describe a substance that is not soluble in another substance, such as oil in water?
Insoluble

A solid substance is dissolved in a liquid. If the solid comes out of solution and settles to the bottom, it is called a precipitate .
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9
6
7
8
5
4
3
1
2
The Periodic Table

K
19
39.093

Atomic Number o Identifies the element o Tells you how many protons an atom has o Tells you how many electrons are contained by a neutral atom of a given element.

Atomic Mass o Average mass of the atom o Equal to number of protons plus number of neutrons . o Electrons have mass BUT the mass is so small we do not factor it in to the overall mass.
You will need to be able to identify the symbol, atomic number, and atomic mass of the first 20 elements on the periodic table.

Fill in the blanks in the table below
Atomic
Number
Name Symbol Atomic Mass

How many protons and neutrons does one atom of the following elements contain?
Hydrogen H
Helium He
1.0079
4.0026
Lithium Li 6.941
Element
Oxygen
Bromine
Protons Neutrons
8
35
8
45
Beryllium Be 9.0122
Boron B 10.811
Carbon-14 6
Atomic Number 53 53
8
74
Carbon
Nitrogen
Oxygen
C
N
O
12.011
14.007
15.999
Fluorine F 18.998
Atomic Number 10 10 10
10 Neon Ne 20.180
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Isotopes

The atomic mass of each atom represents an average of all of the individual isotopes of that element.

Two atoms contain the same number of protons but different numbers of neutrons .

Isotopes are atoms of the same element , but have different masses .

Isotopes with an unstable nucleus will tend to breakdown or decay; these atoms are called radioactive and will release energy in the form of nuclear radiation as they decay.
Metals, non metals, and metalloids

Label the table to indicate the location of metals, non metals and metalloids.

The horizontal rows on the periodic table are called periods.

The vertical columns on the period table are called families or groups .
Oxidation States

In order to become stable, atoms will gain or lose a certain number of electrons .

The goal is to have a full outer shell (octet rule).

A full outer shell usually contains eight electrons.

When atoms gain or lose electrons, they become ions and take on a certain charge. o This charge is referred to as the oxidation number.

On the table above, label the oxidation number for each group.
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Important groups on the periodic table

Alkali Metals o Group 1 o 1 valance electron o Oxidation Number = +1 o Highly reactive

Alkaline Earth Metals o Group 2 o 2 valance electrons o Oxidation Number = +2 o Harder, Denser, Stronger than Alkali
Metals o Very reactive, but less reactive than
Alkali Metals


Transition o Groups 3-12 o Varied oxidation numbers o Not as reactive as groups 1 and 2.
Halogens o o
Metals
Group 17
7 valance electrons o Oxidation Number = -1 o Most reactive non-metals o Combine with metals o NaCl, KBr, MgBr
Bonding

When forming compounds, atoms will bond in a way that leads to an overall charge of zero .

Bonding is due to interactions of the electron clouds that surround an atom.

Types of bonds o ionic o covalent
Ionic Bonds

Formed between a metal and a non-metal .

Forms a compound—not a molecule.

Involves gain/loss of electrons.

Produces compound with net charge of zero.

How to predict bonding pattern: o Na + Cl NaCl o Ca + Br o Ba + I o Mg + O o Al + O
CaBr
BaI
MgO
Al
2
2
O
3
2
Covalent Bonds

Involves the sharing of electrons.

Produces a molecule.

Formed between two non-metals

Examples o Water (H
2
O) o Glucose (C
6
H
12
O
6
) o Hydrogen gas (H
2
)

Diatomic molecules: H
2
, F
2
, Cl
2
, Br
2
, I
2
, N
2
, O
2

Noble Gases o Group 18 o 8 outer electrons; except Helium o will not gain or lose electrons o no oxidation number o Very stable
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Bonding Practice

What type of bond is produced when electrons are shared between atoms? covalent

What type of bond is produced when atoms with opposite charges are attracted to each other? ionic

What type of bond will be produced when the following atoms combine? o C + O covalent o Mg + Cl ionic o O + O covalent o Ba + Br ionic
Periodic Properties

Electron Affinity o The ability of an atom to attract and hold extra electrons.

Electronegativity o The tendency of an atom to attract electrons to itself when combined with another atom.

Ionization energy o Amount of energy required to remove an electron from an atom or ion.

Atomic Radius o One half the distance between two nuclei of like atoms. o A measure of the size of an atom. o What effect does atomic radius have on electron affinity and ionization energy?

Reactivity o Metals

Increases as you move down a family.

Decreases as you move across a period.

Francium is most reactive metal. o Nonmetals

Decreases as you move down a family.

Increases as you move across a period.

Fluorine is the most reactive nonmetal.
1 18
2 13 14 15 16 17
3 4 5 6 7 8 9 10 11 12
Electronegativity Increases
Atomic Radius Decreases

Practice

On the Periodic Table above, identify the following: o o
Metals
Nonmetals

Alkaline Earth Metals

Transition Metals

Halogens

Noble (Inert) Gasses o Alkali Metals

Explain the difference between a family/group and a period on the Periodic Table.

Next to the vertical arrow beside the Periodic Table above, describe what happens to Electronegativity,
Atomic Radius, and Ionization energy as you move down the families.

Below the horizontal arrow under the Periodic Table above, describe what happens to Electronegativity,
Atomic Radius, and Ionization energy as you move across the periods.

List the following elements from highest to lowest electronegativity: o Al, Ca, Cl Cl, Al, Ca o N, Bi, As N, As, Bi o I, Xe, Rb Xe, I, Rb o Cs, Li, K Li, K, Cs

List the following elements from largest to smallest atomic radius: o Al, Ca, Cl Ca, Al, Cl o N, Bi, As Bi, As, N o I, Xe, Rb Rb, I, Xe o Cs, Li, K Cs, K, Li

List the following elements from highest to lowest ionization energy: o Al, Ca, Cl Cl, Al, Ca o I, Xe, Rb Xe, I, Rb o o
N, Bi, As
Cs, Li, K
N, As, Bi
Li, K, Cs
Chemical Reactions
The process by which the atoms of one or more substances are rearranged to form different substances o Reactant
 The starting substance in a chemical reaction. o Product

The substance formed during a chemical reaction. o Catalyst

A substance that increases the rate of a chemical reaction by lowering activation energies but is not itself consumed in the reaction. o Chemical Equation

A statement using chemical formulas to describe the identities and relative amounts of the reactants and products involved in the chemical reaction. o Law of Conservation of Matter
 Matter is neither created nor destroyed

All chemical reactions should be balanced ; the mass of the products should equal the mass of the reactants. o In the chemical reaction below, label the components that are circled.
2
4

2
4
2
Subscript
Yield Sign
Coefficient
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
Types of reactions o Synthesis

Two or more substances react to yield a single product.

2H
2
+ O
2

2H
2
O o Decomposition

A single compound breaks down into two or more elements or compounds.
 2H
2
O  2H
2
+ O
2 o Single Displacement/Replacement

The atoms of one element replace the atoms of another element in a compound.
 2AgNO
3
+ Cu  Cu(NO
3
)
2
+ 2Ag o Double Displacement/Replacement

Involves the exchange of positive ions between two compounds.

AgNO
3
+ KCl

AgCl( s ) + KNO
3 o Combustion

Occurs when a substance reacts with oxygen, releasing _______ in the form of heat and light.

CH
4
+ 2O
2

2H
2
O + CO
2 o Dehydration

Occurs when monomers combine with the loss of a water molecule.

C
6
H
12
O
6
+ C
6
H
12
O
6

C
12
H
22
O
11
+ H
2
O o Exothermic Reaction: Energy is released o Endothermic Reaction: Energy is absorbed
Practice
Identify each reaction below
1.
2C
3
H
7
OH + 9O
2

6CO
2
+ 8H
2
O
2.
Ca
3
(PO
4
)
2
+ 3H
2
SO
4

3CaSO
4
+ 2H
3
PO
4
3.
4.
H
C
2
3
O + SO
H
8
+ 5O
3
2


H
2
SO
4
4 H
2
O + 3CO
2
5.
2KClO
3

2KCl + 3O
2
6.
2KI + Cl
2

2KCl + I
2 combustion double replacement synthesis combustion decomposition single replacement
Chemical and Physical Changes

Chemical change o A change in the arrangement of atoms . o A change where you end up with a new and different substance from which you started. o Combustion, Fermentation, Electrolysis, Rusting/Oxidation, Tarnishing, Souring of Milk, “chemical reactions” o Examples

2H
2
O

2H
2
+ O
2

C
6
H
12
O
6
+ 6O
2

6CO
2
+ 6H
2
O

HCl + NaOH

NaCl + H
2
O

Physical Change o A change in a physical property of a substance. o End up with same substance as original. o Phase changes

H
2
O(s)

H
2
O(l)

H
2
O(g) o Dissolving, Melting, Freezing o Breaking into smaller particles
- 38 -

Practice

Classify each of the following as a chemical or a physical change:
1.
boiling water
2.
bleaching clothes physical change chemical change
3.
drying clothes
4.
slicing potatoes physical physical change change
5.
making coffee
6.
silver tarnishing physical change chemical change
7.
cooking a hamburger chemical change
8.
making kool aid physical change
Acids and Bases

Acid o Donates H + when dissolved in water. o Acidic solutions have more H
+
than OH
-
. o pH less than 7 o Examples

HCl

Lemon juice

Vinegar

H
2
SO
4

Stomach Acid

Base o Donates OH when dissolved in water. o Basic solutions have more OH than H + . o pH greater than 7 o Examples
 NaOH

NH
3
(ammonia)

How is ammonia a base if it does not have OH
-
?

Acid and Base Terms o Neutralization: an acid reacts with a base to produce a neutral solution.

Produces a salt and water .

HCl + NaOH

NaCl + H
2
O o Hydrogen ion: H
+ o Hydroxide ion: OH
o Indicator: a compound that changes color in the presence of an acid or base.

Phenolphthalein

Litmus paper: red (acid), blue (base) o pH: a measure of the hydrogen ion concentration in a solution.

Acid Rain o Normal Rain is slightly acidic due to dissolved CO
2 o Pollutants such as sulfur oxides and nitrogen oxides decrease the pH further. o Rain with a pH less than 5.5
is considered acid rain. o How would acid rain affect plants? o How would acid rain affect buildings and monuments?
- 39 -

Energy Transformations
States of Matter

Matter exists in three primary states o Solid o Liquid o Gas

Solid o Particles closest together o Most dense o Definite shape and volume o Strongest intermolecular forces o Least amount of particle motion (kinetic energy)

Liquid o Particles further apart o Particles have greater range of motion compared to solid o Less dense o Definite volume, but not definite shape o Takes the shape of its container

Gas o Weaker intermolecular forces o Particles farthest apart o Greater particle motion and energy content than solids and liquids. o Least dense o No definite shape or volume; compressible o Takes the shape of its container o Weakest intermolecular forces o Random collisions between particles
Conversion between States
Melting
Evaporation/Vaporization
Condensation
Phase Changes

Melting o Solid
 liquid

Vaporization/Evaporation (Boiling) o liquid
 gas

Freezing o liquid
 solid

Condensation o gas
 liquid

Sublimation o solid
 gas
Freezing
Phase changes require a gain or loss of energy.
- 40 -

Thermodynamics
Thermodynamics
 “Movement of
Heat
”

The study of heat and its transformation to mechanical energy.
Temperature

Tells us how warm or cold an object is relative to some standard.

A measure of the average kinetic energy of a substance.

Temperature is measured using a thermometer .
Temperature Scales

Celsius (
0
C)

Fahrenheit (
0
F)

Kelvin (K)
Important Temperatures
Scale Absolute Zero Freezing Point of Water Boiling Point of Water
Kelvin
Celsius
0K
0 0 C
32 0 F
100 0 C
212 0 F Fahrenheit
What Causes Temperature?

Kinetic-Molecular Theory o Matter made up of tiny particles that are always in motion . o As the particles gain energy, they move faster . o Faster moving particles have greater average kinetic energy . o The more kinetic energy particles have, the greater the temperature of the object or substance.
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