Review for Exam 1

CH 1-2 Concepts to know Classification of matter: pure substances & mixtures

Homogeneous vs Heterogeneous Distinguish the difference between chemical and physical

properties & changes We represent uncertainty with significant figures

You do not need to memorize Sig Fig rules Scientific Notation Conversions within the metric system and non metric units

Temperature conversions Density & Specific gravity Familiarity with how compounds will be drawn

Molecular formulas Structure of an atom: protons, neutrons, electrons

Atomic number, isotope mass number, atomic weight Navigate the periodic table: properties shared within a

group, trends, metals/metalloids/nonmetals Determine valance electrons & draw electron dot

representations Ionization Energy & Atomic Size

Conversions & Equations To Memorize

Unit Conversions

For metric units (m, kg, s, K, mole):mega (M) 106

kilo (k) 103

centi (c) 10-2

milli (m) 10-3

micro (μ) 10-6

nano (n) 10-9

Pico (p) 10-12

Time conversions: dhrms

1 mL = 1 cm3

T(kelvin) = T(°C) + 273.15

Equations

Density = mass / Volumed = m/V

dH2O = 1 g/mL = 1 g/cm3

Specific Gravity = density substance / density of water

Coefficient:A number between1 and 10

y x 10x

Exponent:Any positive or negativewhole number

Elements & Molecules

X = Element symbol (ie O = oxygen)A = Isotope Mass Number = # protons + # neutronsZ = Atomic Number = # protons

6

C12.01

atomic number

element symbol

atomic weight (amu) = weighted average of atomic weight of isotopes

Elements on the Periodic

Table

Molecular Formula: AxBy Ex: CH3O

H C H

H

H

Drawing Molecules:

Methane CH4

Properties of Metals, Nonmetals, Metalloids

Metals Nonmetals Metalloids

• Metallic luster, malleable, ductile, hardness variable

• Conduct heat and electricity

• Solids at room temperature with the exception of Hg

• Chemical reactivity varies greatly: Au, Pt unreactive while Na, K very reactive

• Brittle, dull

• Insulators, non-conductors of electricity and heat

• Chemical reactivity varies

• Exist mostly as compounds rather then pure elements

• Many are gases, some are solids at room temp, only Br2 is a liquid.

• Properties intermediate between metals and nonmetals

• Metallic shine but brittle

• Semiconductors: conduct electricity but not as well as metals: examples are silicon and germanium

Valence Electrons

Example: Determine the valence electrons of Selenium (Se):1. Find Se on the periodic table2. Focus on just the column Se is in3. Column number indicates number of e-

SeElectron Dot Symbols:

Represent the valence electrons by drawing them around the element symbol for Selenium.

Periodic Trends

Ionization EnergyINCREASING

SizeINCREASING

CHAPTER 3-4: Concepts to Know

The difference between ionic and covalent bonds Define cations and anions Predict cation/anion charge using the octet rule or group number

Familiar with metals with multiple potential charges (do not need to memorize)

Determine ionic compound formulas from the name of a compound or from the elements that compose it. Criss-cross rule

Naming ionic compounds and covalent molecules Familiar with polyatomic ions (do not need to memorize but must

be able to recognize) Draw lewis dot structures Determine molecular geometry Identify polar bonds Determine dipole moment of molecules

Need to Memorize

Ionic vs Covalent Bonding

Ionic Bonds result from electrostatic attraction between a cation and anion: metal-nonmetal (with the

exception of NH4+ and H3O+ cations).

Covalent bonds result from the sharing of electronsbetween two atoms: nonmetal-nonmetal.

Li FIonic Bonds Covalent Bonds

Naming

HOW TO Name an Ionic Compound

Step [1]

Determine the charge on the cation.

Step [2]

Name the cation and the anion If the cation could be

multiple charges indicate the charge with roman numerals or with a –ous / -ic suffix.

Step [3]

Write the name of the cation first then the name of the anion

HOW TO Name a Covalent Molecule

Step [1]

Name the first nonmetal by its element name and the second using the suffix“-ide.”

Step [2]

Add prefixes to show the number of atoms of each element.

Predicting Cations & Anions

the anion charge = 8 – group number

the cation charge = the group number

Octet Rule

only 6 e− on B

B

F

FF

10 e− on P 12 e− on S

S

O

OHHO

O

P

O

OHHO

OH

The octet rule: a main group element is especially stable when it possesses an octet of e− in its outer shell.

octet = 8 valence e−

Exceptions (need to memorize):

Ionic Compound Formulas

HOW TO Write a Formula for an Ionic Compound

Step [1]

Identify which element is the cationand which is the anion.

Step [2]

Determine how many of each ion type is needed for an overall charge of zero. When the cation and anion have different

charges, use the ion charges to determine the number of ions of each needed.

Step [3]

To write the formula, place the cationfirst and then the anion, and omit charges.

Lewis Dot Structures

Step [1]

Step [2]

Step [3]

Arrange the atoms next to each other that you think are bonded together. Place H and halogens on the periphery, since they can only form one bond.

Count the valence electrons. The sum gives the total number of e− that must be used in the Lewis structure. For each atom the number of bonds = 8 – valence electrons.

Arrange the electrons around the atoms. Place one bond (two e−) between every two atoms. Use all remaining electrons to fill octets with lone pairs, beginning with atoms on the periphery.

NH3

N HH H

Nitrogen has 5 valence electrons, so it will have 8 – 5 = 3 bonds.

Hydrogen will have 2-1 = 1 bond.There are 8 total valance electrons

H NH

H1 lone pair:

2

3 bonds:

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Total e-

8

= total valence e-

Resonance StructuresResonance structures exist when there are multiple lewis dot structures with different electron arrangements with the same connectivity between atoms. Resonance structures help us understand delocalization (spreading) of charge within a molecule that stabilizes the anion or cation.

Other Examples: CO32- and O3

Molecular Shape

Periodic Trend: Electronegativity

ElectronegativityINCREASING

Polarity

1. Assess the relative electronegativity of atoms bonded together, if there is a difference it is a polar bond.

2. Indicate polar bonds with δ+ / δ - or3. If polarity of bonds does not cancel draw the overall

dipole moment of the molecule using

Electron density is disproportionately distributed over the molecule. Above red indicates partial negative charge, or greater electron density, and blue indicates partial positive charge.Effectively oxygen is hogging the electrons

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CH 5 Concepts to know

Smith. General Organic & Biolocial Chemistry 2nd Ed.

Define a chemical reaction

Correctly write a chemical reaction

Balance reactions by inspection

Calculate molecular mass for any compound or molecule

Apply mole ratios within molecules and between molecules.

Solve stoichiometry problems

Convert between mass and moles

Identify limiting reagent

Calculate percent yield

Identify reduction and oxidation equations and pick out the

compound being reduced or oxidized

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Need to Memorize

Smith. General Organic & Biolocial Chemistry 2nd Ed.

6.02 x 1023 is Avogadro’s number.

Oxidation is the loss of electrons from an atom.Reducing agents are oxidized

Reduction is the gain of electrons by an atom.Oxidizing agents are reduced.

Smith. General Organic & Biolocial Chemistry 2nd Ed. 22

aA (physical state) + bB (state) cC (state) + dD (state)

Writing and Balancing Equations

HOW TO Balance a Chemical Equation

Step [1]

Write the equation with the correct formulas.•The subscripts in a formula can never be changed to balance an equation, because changing a subscript changes the identity of a compound.

Step [2]

Balance the equation with coefficients oneelement at a time.

Step [3]

Check to make sure that the smallest setof whole numbers is used.

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aA + bB cC + dDmass A mass B mass C mass D

moles A moles B moles C moles Da:b b:c c:d

a:c

a:d

FW/MM

Solve Stoichiometry Problems

mol mol

g g

FW/MM FW/MM FW/MM

Note:FW/MM meansFormula wt. orMolar mass

Limiting Reactant

Compare the actual amount of each reactant to the amount required in the balanced equation to determine how many times the “reaction can be run”

Use the amount of the limiting reactant to calculate how much product can be produced

Smith. General Organic & Biolocial Chemistry 2nd Ed. 25

Zn + Cu2+ Zn2+ + Cu

Zn loses 2 e–

Cu2+ gains 2 e−

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Oxidation half reaction: Zn Zn2+ + 2 e−

Each of these processes can be written as an individual half reaction:

loss of e−

Reduction half reaction: Cu2+ + 2e− Cu

gain of e−

Redox Half Reactions