Unit 1 – Matter and Measure Unit 2 – Atomic Theory and Structure Unit 3 – Nuclear Chemistry Unit 4 – Periodic Table Unit 5 - Bonding Unit 6 – Naming and Moles

Regents Chemistry Midterm Review

Matter and MeasureChapters 1-3

What is Chemistry?◦ Study of matter and the changes it undergoes

Branches◦ Organic◦ Physical◦ Analytical◦ Biochemical◦ Inorganic

Chemistry

Anything that has mass and takes up space, volume

Classified into two categories◦ Substances (Pure)◦ Mixtures

Matter

Matter

Element◦ simplest form of matter that has a unique set of

properties.◦ Can’t be broken down by chemical means

Compounds◦ substance of two or more elements chemically

combined in a fixed proportion◦ Can be broken down by chemical means

Pure Substances

Physical blend of two or more substances

Two Types:◦ Homogeneous

Composition is uniform throughout

◦ Heterogeneous Composition is not uniform throughout

Mixtures

Differences in physical properties can be used to separate mixtures◦ Filtration – Separates solids from liquids in

heterogeneous mixtures◦ Distillation – Separates homogeneous liquid

mixtures based on different boiling points◦ Evaporation – evaporate away liquid to leave

solid◦ Chromatography – separation of substances

based on polarity

Separating Mixtures

Solid (s)◦ Definite shape and volume◦ Particles are packed tightly together in a regular

geometric pattern Liquid (l)

◦ Definite volume, takes shape of container◦ Particles can slide past each other

Gas (g)◦ Takes shape and volume of container◦ Particles are spread very far apart

Phases(States) of Matter

Solid Liquid Melting Liquid Solid Freezing Liquid Gas Vaporization Gas Liquid Condensation Solid Gas Sublimation Gas Solid Deposition

Temperature does NOT change during a phase change

Phase Changes

Physical Property◦ quality or condition of a substance that can be

observed or measured without changing the substance’s composition

◦ Ex: Color, shape, size, mass

Physical Change◦ some properties change, but the composition

remains the same◦ Can be reversible or irreversible◦ Ex: melting, freezing, tearing

Identifying Substances

Chemical Change◦ change that produces matter with a different

composition than the original matter◦ Ex. burning, rusting, decomposing, exploding,

corroding

Chemical property◦ property that can only be observed by changing

the composition of the substance.◦ Ex: Reactivity with acids, reactivity with oxygen

Identifying Substances (cont)

Exothermic◦ Process when energy is released or given off◦ Ex: Burning, freezing

Endothermic◦ Process when energy is absorbed or taken in◦ Ex: Melting

Energy Exchanges

Observation◦ using five senses to make observations.

Hypothesis◦ proposed explanation for an observation.

Experiment◦ procedure used to test a hypothesis.

Analyze Data◦ check to see if results support hypothesis.

Theory◦ well tested explanation for a broad set of observations.

Law◦ concise statement that summarizes the results of many

observations and experiments.

Scientific Method

Shorthand◦ If the decimal point is present, start counting

digits from the LEFT side, starting with the first non-zero digit.

Significant Figures

1 2 3

0.00310

(3 sig. figs.)

Shorthand ◦ If the decimal point is absent, start counting

digits from the RIGHT side, starting with the first non-zero digit.

Significant Figures (cont)

3 2 1

31,400

(3 sig. figs.)

Addition and Subtraction◦ Answer has to have the same number of decimal

places as least decimal places in what you are adding or subtracting

Multiplication and Division◦ Answer has to have same number of Sigfigs as

least number of Sigfigs in what you are multiplying or dividing

SigFigs for Math

Percent Error

Measure related to the heat of an object

Measured in °Celsius or Kelvin(no degrees)

Conversion

Temperature

273 CK

Amount of matter in a given amount of space

Amount of mass in a given volume

Density

V

mD

Back

Atomic Theory and Structure

Chapters 4-5

Atom Atoms are made of subatomic particles

◦ Protons◦ Neutrons◦ Electrons

Electron Discovered first Negative charge (-1) Approx mass ~ 0u Found outside of nucleus

Valence Electron◦ Electrons in the outermost energy level

Proton Discovered second Positive charge (+1) Approx mass ~ 1u Found inside nucleus

Neutron Discovered last No charge (0) Approx mass ~ 1 atomic mass unit (u)

◦ Just slightly larger than a proton Found inside nucleus

Atomic Structure Atoms have no net charge

◦ # of electrons = # of protons

Nucleus◦ Center of atom, contains protons and neutrons◦ Positive charge

Atomic Structure Atomic Number

◦ Number of protons◦ All atoms of the same element have the same

number of protons

Mass Number◦ Number of protons and neutrons in an atom◦ # of Neutrons = Mass Number – Atomic Number

Chemical Symbols

Cl-35 Chlorine-35

Cl3517

Mass Number

Atomic Number

Atomic Structure Isotope

◦ atoms of the same element with different number of neutrons

Ion◦ Atom or group of atoms that have gained or lost

one or more electrons◦ Have a charge

Average Atomic Mass Atomic Mass

◦ Weighted average based on the relative abundance and mass number for all naturally occurring isotopes

Relative Abundance◦ Percent of each naturally occurring isotope found

in nature

Atomic Mass C-12 98.9% C-13 1.1%

Carbon = 0.989*12 + 0.011*13 = 12.011u

Atomic Theories Dalton’s Atomic Model

◦ Also called Hard Sphere Model◦ First model

Plum Pudding Model◦ Uniform positive sphere with negatively charged

electrons embedded within.◦ Came as a result of discovery of electron

Rutherford Gold Foil Experiment Shot alpha particles at gold foil Most went through, some were deflected

back

Conclusions◦ Atom is Mostly Empty Space◦ Dense positive core (nucleus)

Atomic Theories Rutherford Model

◦ Dense positive core (nucleus)◦ Electrons moving randomly around nucleus

Bohr Model◦ Dense positive core (nucleus)◦ Electrons in specified circular paths, called energy

levels

Atomic Theories Wave Mechanical Model

◦ Dense positive core (nucleus)◦ Electrons in orbitals

Regions of space where there is a high probability of finding an electron

◦ Modern (current) Model AKA Quantum Mechanical Model, Electron Cloud

Model

Electron Configuration The way in which electrons are arranged in

the atom

Ground State◦ When the electrons are in the lowest available

energy level (shown on reference tables)

Excited State◦ When one or more electrons are not in the lowest

available energy level

Bohr Model Each energy level can only hold up to a

certain number of electrons

Level 1 2 electrons Level 2 8 electrons Level 3 18 electrons Level 4 32 electrons

Valence Electrons

Electrons in the outermost energy level

Energy Level Transitions Electrons can move between energy levels

Gaining energy will move an electron outward to a higher energy level

When an electron falls inward to a lower energy level, it releases a certain amount of energy as light

Back

Nuclear ChemistryChapter 25

Radioisotopes Nuclei of unstable isotopes are called

radioisotopes.

An unstable nucleus releases energy by emitting radiation during the process of radioactive decay◦ Mass and/or energy

Radiation Late 1800’s – discovery of radiation

Three Types◦ Alpha◦ Beta◦ Gamma

Symbols

He42

42

e0101

Alpha

Gamma

Beta

00

Radiation

Radiation

Radiation Three Types

RadiationWhat it

resembles Mass Charge Strength

AlphaHelium Nucleus

4 +2 Weakest

Beta Electron 0 -1 Middle

GammaLight wave

0 0Stronges

t

Nuclear Stability For smaller atoms a ratio of 1:1 neutrons to

protons helps to maintain stability◦ C-12, N-14, O-16

For larger atoms, more neutrons than protons are required to maintain stability◦ Pb-207, Au-198, Ta-181

Transmutations Any reaction where one element is

transformed into a different element◦ Nuclear Reactions

Natural◦ Has one reactant◦ Alpha and Beta Decay

Artificial◦ Has more than one reactant◦ Particle Accelerators

Radioactive Decay Radioisotopes will undergo decay reactions

to become more stable

Alpha Decay

Beta DecayAtFr 216

8542

22087

YSr 9039

01

9038

Fission Splitting of a larger atom into two or more

smaller pieces◦ Nuclear Power Plants

One Example:

nKrBanU 10

9236

14156

10

23592 3

Fusion Joining of two or more smaller pieces to

make a larger piece◦ Sun, Stars

Examples: 01

42

11 24 HeH

nHeHH 10

32

21

21

nHeHH 10

42

31

21

Energy Production Energy is produced by a small amount of

mass being converted to energy◦ More energy is produced by fusion than any other

source

E=mc2

Fission vs. Fusion Advantages of Fission

◦ Produces a lot of energy◦ Can be a controlled reaction◦ Material is somewhat abundant

Fission vs. Fusion Disadvantages of Fission

◦ Uses hazardous material◦ Produces hazardous material

Long Half Life◦ Reaction can run out of control.◦ Limited amount of fissionable material

Fission vs. Fusion Advantages of Fusion

◦ Lighter weight material◦ Easily available material◦ Produces waste that is lighter and has shorter

half-life◦ Produces more energy than fission

Fission vs. Fusion Disadvantages of Fusion

◦ Must be done at very high temperatures Only been able to attain 3,000,000K

◦ Have not been able to sustain stable reaction for energy production

Half Life Amount of time for half of a sample to

decay into a new element

Parent Atoms◦ Undecayed atoms

Daughter Atoms◦ Decayed atoms

Half Life

Number of Half-lives

Fraction left

0 1

1 1/2

2 1/4

3 1/8

4 1/16

5 1/32

Half Life Number of half-lives

T

tHalfLives #

t = amount of time elapsed

T = half-life

Half Life Fraction Remaining

T

t

maining

Fraction

2

1

Re

t = amount of time elapsed

T = half-life

Fraction Remaining

T

t

2

1Mass Left

Original Mass=

Example How many half lives does it take for a

sample of C-14 to be 11430 yrs old?

25715

11430

y

y

T

t

Example What fraction of P-32 is left after

42.84days?

8

1

2

1

2

1

2

13

28.14

84.42

d

d

T

t

Example How long will a sample of Rn-222 take to

decay down to 1/4 of the original sample?

2823.3

2

1

2

1

4

1

d

X2

823.3

d

x

7.646d

Practice How much Carbon-14 was originally in a

sample that contains 4g of C-14 and is 17145 years old?

8

1

2

1

2

143

5715

17145

y

y

x

g

32g

More Practice How much 226Ra will be left in a sample that

is 4797 years old, if it initially contained 408g?

8

1

2

1

2

1

408

31599

4797

y

y

g

x

51g

And One More…. What is the half life of a sample that started

with 144g and has only 9g left after 28days?

428

2

1

16

1

2

1

144

9

x

d

g

g4

28

x

d

7d

Radioisotopes You must know these radioisotopes and

uses◦ I-131

Diagnosing and treating thyroid disorders◦ Co-60

Treating cancer

Radioisotopes You must know these radioisotopes and

uses◦ C-14

Dating once-living organisms Compare to C-12

◦ U-238 Dating geologic formations Compare to Pb-206

Back

Periodic TableChapter 6

Dmitri Mendeleev 1869 - Russian chemist and teacher,

proposed a table for organizing elements◦ Arranged the elements by increasing atomic

mass.◦ Left spaces for elements not yet discovered◦ Predicted very closely the properties of Ge, Ga,

Sc, and 5 others

Periodic Table Arranged in order of increasing atomic

number◦ Columns are called Groups

Numbered 1-18◦ Rows are called Periods

Periodic Law Periodic Law – When elements are arranged

in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

Group Names Group 1 - Alkali Metals Group 2 - Alkaline earth metals Group 17 – Halogens Group 18 - Inert or Noble gases Groups 3-11 – Transition Metals Bottom 2 rows – Inner Transition

Valence Electrons Electrons in outermost occupied energy

level

Valence Electrons are responsible for most chemical properties◦ Elements in the same group have similar

properties because they have the same number of valence electrons

Phases at STP Most elements are solids at STP

Hg and Br are liquids at STP

H, N, O, F, Cl and Noble Gases are all gases at STP

Classifying Elements Elements are classified into 3 groups based

on their properties:

Metals – Left and Middle

Nonmetals – Right

Metalloids - Staircase

Metals Good conductors of heat and electrical

current High luster or sheen Many are ductile, meaning they can be

drawn into wires Most are malleable, meaning they can be

hammered into thin sheets

Nonmetals Most are gases at room temperature, some

are solids, and one is liquid

Most are poor conductors

Most solids are brittle

Metalloids B, Si, Ge, As, Sb, Te

Have properties of both metals and nonmetals, based on conditions

Exceptions:◦ Al and Po are metals◦ At is a nonmetal

Diatomic Elements Eight elements are diatomic molecules

when alone in nature (exist as two atoms bonded together)

H2, N2, O2, F2, Cl2, Br2, I2, At2

Hydrogen and the Magic 7

Ions Atom, or group of atoms, that has gained or

lost electrons

Cation – positive ion Anion – negative ion

Ions When an atom loses an electron, it becomes

positively charged◦ The radius becomes smaller◦ Metals tend to lose electrons

When an atom gains an electron, it becomes negatively charged◦ The radius becomes larger◦ Nonmetals tend to gain electrons

Properties (Table S) Atomic Radius

◦ Size of the atom

Ionic Radius◦ Size of an ion

Properties (Table S) First Ionization Energy

◦ Amount of energy required to remove the outermost electron

Electronegativity◦ Ability of an atom to attract an electron from

another atom when in a compound. Noble gases are omitted, don’t form compounds

Periodic Table Trends Atomic Number

◦ increases across a period.◦ increases down a group

Atomic Mass ◦ generally increases across a period.◦ increases down a group.

Periodic Table Trends Atomic Radius

◦ Decreases across a period◦ Increases down a group

Ionic Radius◦ Decreases for positive/negative ions across a

period◦ Increases down a group

Periodic Table Trends First Ionization Energy

◦ Tends to increase across a period◦ Tends to decrease down a group

Electronegativity ◦ Tends to increase across a period◦ Tends to decrease down a group

Metallic/Nonmetallic Character Metallic Character increases as you move

towards the lower left◦ Most Metallic Element is Francium, Fr

Non-Metallic Character increases as you move towards upper right◦ Most nonmetallic element is Fluorine, F

Trends Summary

Property Period (LR) Group (TB)

Atomic Number

Atomic Mass

Atomic Radius

Ionic Radius

Ionization Energy

Electronegativity

Reactivity Elements that are more reactive tend to

either gain or lose electrons very easily Elements that lose electrons easily have low

IE and low EN◦ Lower left, Fr

Elements that gain electrons easily have high IE and high EN◦ Upper right, F

Back

BondingChapters 7-8

Octet Rule Atoms tend to lose or gain electrons to

achieve a full valence shell (8)◦ Exception: First Energy Level is full with 2

electrons

Electron Dot Structures Diagrams that show valence electrons,

usually as dots◦ AKA Lewis Electron Dot Diagrams

Rules◦ Start on any side◦ First two get paired together◦ Next three are separated◦ Fill in as needed

O

Ions Atoms that have gained or lost electrons,

and now have a charge

Must show charge

Na+ F- O-2

Compounds Two Main Types of Compounds

◦ Ionic◦ Molecular (Covalent)

Based on type of bonding involved

Bonding Bond

◦ Shared or exchanged electrons that hold two atoms together

Three Main Types◦ Covalent◦ Ionic◦ Metallic

Covalent Bonds Electrons are shared between two atoms to

hold them together◦ Each atom will try to achieve a full valence shell◦ 2 nonmetals

Two types of covalent bonds◦ Non-Polar Covalent – Shared equally◦ Polar Covalent – Shared unequally

Covalent Bonding

H2

O2

N2

HHSingle Bond

Double Bond

OO

NNTriple Bond

Covalent Bonding

H2O

CO2

OH

HSingleBond

SingleBond

CO O

More Examples HCl

NH3

CH4

ClH

NHH

H

CHH

HH

Bonding Ionic Bond

◦ Electrons are transferred from one atom to another (one gives, one takes)

◦ Metal and nonmetal, NaCl Large electronegativity difference

◦ Polyatomic ion, Mg(NO3)2

More than 2 elements

Properties Ionic Compounds

◦ Most ionic compounds are hard, crystalline solids at room temperature

◦ High melting points

◦ Mostly soluble in water

◦ Can conduct an electric current when melted or dissolved in water(aq).

Properties Covalent Compounds

◦ Most molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.

Determining Bond Type Bond type is based on electronegativity

difference (ΔEN) between two bonding atoms

Nonpolar Covalent Bond2 of the same nonmetals (no difference in electronegativity)

Polar Covalent Bond2 different nonmetals (small difference in electronegativity)

Ionic BondMetal and a nonmetal (large difference in electronegativity)

Determining Bond Polarity The larger the difference in

electronegativity, the more polar the bond. Which is more polar?

H IH BrH ClH F 1.8

1.0

0.8

0.5

ΔEN

BiggestMost Polar

Dot Structures Shows valence electrons Must show charge for Ions

NaCl

Na+

Cl-

Polyatomics

Compounds with polyatomic ions contain BOTH ionic and covalent bonds◦ Example: NaNO3

Na+

-NOO

O

Metallic Bonding Bonding within metallic samples is due to

highly mobile valence electrons◦ Free flowing valence electrons◦ “Sea of Electrons”

Allotropes Two or more different molecular forms of

the same element in the same physical state◦ Different properties because they have different

molecular structures◦ O2 vs O3

◦ Diamond, Graphite, Fullerenes (pictured on next slide)

Allotropes

Network Solids All atoms in a network solid are covalently

bonded together◦ very high melting and boiling points

Examples◦ Diamonds ( C )◦ Graphite ( C )◦ Silicon Dioxide (SiO2)◦ Silicon Carbide (SiC)

Bond Energy When two atoms form a bond, energy is

released◦ Example: Cl + Cl Cl2 + energy

Energy needs to be added to break a bond◦ Example: Cl2 + energy Cl + Cl

Molecular Polarity Polar Molecule

◦ one end of a molecule is slightly negative(δ-) and the other end is slightly positive(δ+).

◦ Asymmetrical charge distribution

Nonpolar Molecule◦ Can not be separated into different ends◦ Symmetrical charge distribution

Polar Molecules

H2O

HCl

NH3

O H

H

δ-

δ+

δ+

H Clδ+ δ-

N

H

δ-

δ+ δ+

δ+

N HH

H

Nonpolar Examples

CH4

CO2

C H

H

H

H

δ+δ-

δ+

δ+

δ+

O=C=Oδ+

δ- δ-

Polarity Ionic Compounds are Ionic

Nonpolar Covalent Bonds always indicate Nonpolar Molecules

Polar Covalent Bonds◦ Determine Symmetry

Polarity Nonpolar Molecules

◦ CH4, CO2, H2, N2, O2, …

Polar Molecules◦ H2O, HCl, HBr, NH3, …

“Like Dissolves Like” Polar and Ionic substances will dissolve in

other Polar Substances

Nonpolar substance will dissolve in other nonpolar substances

Intermolecular Forces Intermolecular Forces of Attraction

◦ attraction between two molecules or ions that hold them together (not a bond)

◦ Determines melting and boiling points of compounds

Stronger intermolecular forces, higher melting and boiling points

Intermolecular Forces Van der Waals

◦ Dispersion◦ Dipole-Dipole

Molecule-Ion Hydrogen Bonding

Weakest

Strongest

Hydrogen Bonding Hydrogen bonded to N, O, or F, is attracted

to the N, O, or F of another molecule. Not actual bond, just attraction

H F H F

Hydrogen “Bond”

Back

Naming and MolesChapter 9-10

Naming Ions Positive Ions, cations, simply retain their

name.◦ Na+ Sodium Ion◦ Mg2+ Magnesium Ion

Negative Ions, anions, change ending of element to –ide◦ Cl- Chloride Ion◦ Br- Bromide Ion

Polyatomic Ions Selected polyatomic ions are on Table E in

the Reference Tables.

Polyatomic ions keep their names in most chemical names

Naming Systems Ionic System

◦ Metals and Nonmetal, more than 2 elements

Stock System (Roman Numerals)◦ Use when the metal element has more than one

positive oxidation number◦ Roman Numeral is the charge of the metal ion

Binary Covalent System (Prefixes)◦ 2 nonmetals (including metalloids)◦ Second element ends in –ide

Naming Ionic Compounds Name positive ion first, then negative ion.

◦ NaCl Sodium chloride◦ Mg(OH)2 Magnesium hydroxide

Example

Fe3(PO4)2

Iron(II) Phosphate

-3

-3

+6 -6

PO4

PO4

Fe

Fe

Fe+2

+2

+2

Roman Numerals

Cation Charge

Roman Numeral

+1 I

+2 II

+3 III

+4 IV

+5 V

+6 VI

+7 VII

+8 VIII

Binary Covalent Example

N2Cl3◦ Dinitrogen Trichloride

CO2

◦ Carbon Dioxide

PCl5◦ Phosphorus Pentachloride

Prefixes

Number of atoms Prefix

1 mono-

2 di-

3 tri-

4 tetra-

5 penta-

6 hexa-

7 hepta-

8 octa-

Avogadro’s Number 6.02 x 1023

Number of representative particles in a mole

1 mol He = 6.02 x 1023 atoms 1 mol H2 = 6.02 x 1023 molecules

Gram Formula Mass Mass of the formula in g/mol Simply add the atomic masses of each

element in the formula together H2O = 1 + 1 + 16 = 18 g/mol Also known as gram atomic mass, gram

molecular mass, molar mass

Mole - Mass Conversion

Example: 96 g of Oxygen gas = ? mol

molmolg

gmol 3

/32

96#

Practice How many moles are there in 506g of

ethanol, C2H6O?

What is the mass of 8 moles of CCl4?

molg

gx

/46

506

molg

xmol

/1548 1232g CCl4

11 mol C2H6O

Molar Volume At STP, 1 mol of any gas occupies 22.4L of

space

Mole Road Map

Percent Composition

%100*Whole

Part

Percent Composition What is the percent composition of oxygen

in H2SO3?

58.5%

100*82

48

Hydrates Compounds that have a specific number of

water molecules attached

◦ Dot means plus (+)◦ gfm = 159.5 + 5(18) = 249.5g/mol

CuSO4·5H2O

Empirical Formula Simplest Whole-Number ratio of atoms in a

compound Examples C6H12O6 CH2O

Molecular Formula is a multiple of the Empirical Formula

Empirical Formula A molecular formula has an empirical

formula of CH2 and a molecular mass of 28 g/mol.

A molecular formula has an empirical formula of CH2 and a molecular mass of 42 g/mol.

C2H4

C3H6