Study Guidechemistry Comprehensive Exam2014

Study GuideChemistry Comprehensive Exam2014

Memorization: see “Stuff to Memorize Sheet” for more (from teacher)

Ions that Form Soluble Compounds
Group 1 ions (Li+, Na+, ect.)
Ammonium
Nitrate
Acetate
Bicarbonate (aka Hydrogen Carbonate)
Chlorate
Perchlorate

Polyatomic ions

/ Nitrate / / Chlorite
/ Hydrochlorite / / Chromate
/ Phosphite / / Dichromate
/ Cyanide / / Chlorate
/ Hydroxide / / Sulfate
/ carbonate / / Bicarbonate*
/ Sulfite / / Permanganate
/ Phosphate / *** / Acetate
/ Perchlorate / / Ammonium

*AKA Hydrogen Carbonate

Lab stuff

Accurate: A measurement that is close to the accepted or correct value.

Precise: Measurements that are close to one another.

Percent Error: Measures accuracy.

Uncertainty: Measures precision.

Absolute Deviation: The difference from the average. Report using

Percent Error:

Uncertainty: ( means average)

Significant figures

Types of Zeros

Sandwich zero: 505
Leading zero: 0.0005
Trailing zero: 5000

Rules

All non-zero digits are significant
All sandwich zeros are significant
Leading zeros are never significant
Trailing zeros are not significant when there is no decimal and are always significant when there is a decimal after of in-between.

Addition and Subtraction: You must round your answer to the least number of decimal places.

Multiplication and Division: You must round you answer to the least number of significant figures.

Physical and chemical changes/properties

Precipitate: A solid formed when two solutions are mixed.

Spectator Ion: An ion that doesn’t participate in a chemical change.

Physical Change: In a physical change molecules stay together, the substance doesn’t change, bonds are never broken, and the change is reversible.

Chemical Change: In a chemical change bonds are broken and reformed, new substances are formed, and the change may not be reversible.

Periodic table basics

Groups

1a/b-Alkali Metals
2a/b- Alkaline Earth Metals
3b-12b- Transition Metals
7a/17b- The Halogen
8a/18b- Noble Gasses
FOR CHARGES SEE PERIODIC TABLE 1

Metals vs. Metalloids vs. Non-metals

Metals: to the left of the stairstep
Non-metals: to the right of the stairstep
Metalloids: Directly touching the stairstep

Ionic compound: Metal and non-metal

Covalent Compound: only non-metals

Periodic table trends and properties (to compare and explain)

Why are orbital diagrams so important? They allow you to predict the ions of elements, including transition metals.

Valence Electrons: The outermost electrons in an atom, that are included in chemical reactions.

Core Electrons: The inner electrons of an atom, that aren’t included in chemical reactions.

Another name for core electrons is shielding electrons.

Shielding electrons “shield” the outermost electrons from the nucleus “pull”.
Called effective nuclear charge ()

Z is the atomic number or number of protons
S is the number of shielding electrons

The higher the the more the valence electrons are attracted to the nucleus.
The higher the the more likely the element will become an anion(negative ion).

Periodic Trends (See Periodic Table 2)

Homework

Why does atomic radius decrease as you move across the periodic table?

As you move from right to left is increasing. The nucleus “pulls” electrons in closer as a result of the higher.

How are the trends of atomic radius and ionization energy related?

They are inversely related, so as ionization energy increases, atomic radius decreases, left to right in a period. In a group, as you go down, ionization energy decreases and atomic radius increase.

What’s the difference between electronegativity and ionization energy?

Ionization energy removes electrons and electronegativity is abou the attraction of electrons.

What does it mean to be isoelectronic? Identify the largest ion and smallest ion of the group: . Explain why.

Isoelectronic means same number of electrons.

Ex) They all have 10 electrons.

Smallest Ion:.

Largest Ion:

The number of protons is decreasing for the same number of electrons.

The first ionization energy of beryllium is 9.322 eV, the second ionization energy is 18.211 eV, and the third ionization energy is 153.893 eV. Explain why the third ionization energy of beryllium is so much higher than the first two.

The first two are valence electrons and the next is a core electron.

Atomic Radius/size: see above notes
Ionic Radius/size: The radius of an ion.

For cations: the ion of an atom is smaller than the atom. Reason: same number of protons, fewer electrons, so more “pull”.
For anions: the ion of an atom is larger than the atom. Reason: same number of protons, more electrons, so less “pull”.

Electronegativity/Electron Affinity: How much an atom attracts an electron.
Ionization Energy (1st): The energy required to remove an electron from an atom.

1st IE: The energy required to remove the 1st or outermost electron.

Metallicity/Metallic Character: How much like a metal an element is.
Reactivity: The ability to gain or lose electrons more easily.

Metals reactivity increase as you go left and down a group.

Reactive Metals: easily lose their electrons; low ionization energy; large atomic radius; low electronegativity.
Low ionization energy: most important characteristic.

Non-metals reactivity increases as you go right and up a group.

Reactive Non-metals: easily gain electrons; high ionization energy; small atomic radius; high electronegativity.
High electronegativity: most important characteristic.

Factors affecting periodic trends:

: effective nuclear charge (only explains trends across a period); every atom in the same group has the same
Distance/Size: The farther from the nucleus, the less “pull” the nucleus can exert. (Columb’s Potential Energy)
Columb’s Potential Energy:

k- Constant -Charges d- distance V- potential energy

Takes into account both charge and distance.
According to Columb’s PE, as distance decreases, potential energy will increase. If potential energy increases then the electron is more attracted to the nucleus.
If q (charge o the nucleus) increases the potential energy increases.

Naming Compounds

Charges:

Zn (zinc) is always 2+
Ag (silver) is always 1+

Steps

1)Identify the type of compound. Then follow specific steps for each type.

Ionic: metal and non-metal (cation and anion); give and take electron. TIP: Does it have a metal?
Covalent: non-metals only; shared electrons. TIP: Not allowed to have a metal, no H in front.
Acid: TIP: H in front, no metals.

Covalent: prefix (no mono-) and name of 1st non-metal + prefix and name of second metal plus –ide.

Prefixes:

1 / Mono / 6 / Hexa
2 / Di / 7 / Hepta
3 / Tri / 8 / Octa
4 / Tetra / 9 / Nona
5 / Penta / 10 / Deca

Ionic:

1)Binary or polyatomic?

Binary: 2 elements only.
Polyatomic: 3+ elements.

Binary: Name of metal + name of non-metal plus –ide.
Polyatomic: name of metal + mane of polyatomic ion.
Transition Metal Compounds: Need a roman numeral in the name except for zinc (Zn) and Silver (Ag). The Roman numeral represents the charge of the metal. Transition metals have variable charges so they need a roman numeral.

Acidic: H in front, because acids donate H+ ions.

HCl: Chloride hydro-chlor-icacid

Hydro-name of anion- ic +acid
TIP: Hydro- only H and 1 element

: ChloriteChlor-ous acid

Name of anion- + acid

:ChlorateChlor-icacid

Name of anion-ic + acid

If you have a polyatomic ion in your acid, the acid’s name doesn’t include hydro.

Writing Chemical Formulas

: represents the atoms in a molecule (structure); subscripts tell us how many of each atom is present; symbols tell us which elements are in the molecule.

Covalent:

1)Write the symbol of the elements

2)Write the subscripts of the elements according to the prefixes

Ionic:

1)Write the symbol of the cation and its charge.

2)Write the symbol of the anion and its charge.

-ide element (except hydroxide and cyanide)
-ite, -ate  anion is polyatomic ion

3)Change the subscripts so the charges add up to zero

Memory Device: Symbol-Charge, Symbol-Charge, cross it, make it pretty
TIPS:

Make sure the charges add to zero
Reduce subscripts if possible
Parentheses around polyatomic ions
Roman numeral is the charge not the subscript

Acidic: H in from all the time

1)Write the symbol of the anion with charge

2)Add enough H+ to make the charge add up to zero

TIPS:

Use the suffix to determine the anion
Hydro+ic just one element
+ic polyatomic ion ending in –ate
+ous polyatomic ion ending in –ite

Percent composition

Converting Moles (mol.) to grams (g):

Molar mass of a single element is the mass found on the periodic table

Empirical Formulas

Empirical Formula: Represents the ratio between the number of each atom found in the compound. The ratio is a mole to mole ratio.

Ex) 1 mol C atoms: 1 mol H atoms: 3 moles of Cl atoms

Empirical formulas don’t always represent the actual structure of the compound. They only tell you the correct ratio.
There are 3 possible starting points for an empirical formul calculation:

(1)% comp (2) Grams (3) Moles  Mole Ratio (subscripts)

(1)  (2) : change percent to grams (assume 100 grams)
(2)  (3) : divide by molar mass
Step 4 (mole ratio): divide all by smallest number of moles
Why assume 100g? Makes conversion simple.
Why divide by the smallest number of moles? Allows you to make a mole ratio where the smallest number is 1.
TIPS:

What if the ratio isn’t a whole number?

Ex) 1.5 mol O *2 = 3

1 mol N *2= 2

Ends in….

.1 or .9 5ound up/down

.2 or .8 multiply by 5

.3 or .7 multiply by 3

.4 or .6 multiply by 5

.5 multiply by 2

Example Problem: Calculate the empirical formula that contains 38.6% N and 63.3% O.

Molecular Compounds: Share the same ratio of atoms, but their actual structures differ. How do we tell the difference between compounds?

Represents the actual structure of the compound but shares the same ratio as the empirical formula.
Easy way to remember the difference between types of formulas:

Empirical: simplest ratio, not structure

Molecular: structure and ratio

How to convert between empirical and molecular formulas:

Multiply subscripts of empirical by the whole number

Example Problem: The compound ethylene glycol is used in antifreeze. IT contains 38.7% C, 9.75% H, and the rest is oxygen. The molecular weight of ethylene glycol is 62.07 g/mol. What is the molecular formula?

Balancing Chemical Equations

Conservation of Mass: Mass can’t be created or destroyed.

Conservation of mass tells you two important things about equations:

1)The amount of each element must be the same on both sides of the equation.

2)The types of elements must also be the same on both sides of the equation.

We balance reactions to satisfy the Law of Conservation of Mass.

What are you allowed to do to balance an equation?

Change the coefficients of compounds

What are you not allowed to do to balance an equation?

Change the subscripts of an element
Change the elements

TIPS:

Treat polyatomic ions as a unit or group
Write water (H2O) as HOH
No fractions in final answers
Reduce coefficients only if you can reduce all of them
For combustion reactions balance in the following order: H C O

Chemical Equations/reactions

5 basic types:

1)Combustion

2)Synthesis

Oxidation- reduction

3)Decomposition

Oxidation- reduction

4)Single Replacement

Oxidation- reduction

5)Double Replacement

Acid-base
Precipitation

Oxidation Reduction (redox), acid-base (a/b), and precipitation (ppt) are special classes of the 5 main types

To predict the products of a reaction always identify the simple type (1 of the 5) first and then worry about the special class.

Combustion (1): be able to write both reactants and products from words.

A reaction where a compound, usually a hydrocarbon (mostly made of carbon), reacts with oxygen gas (O2) to form carbon dioxide (CO2) and water (H2O)
Ex) CH4, methane, combusts. Write the reaction.

Synthesis (2): be able to identify the type; no prediction

A reaction where simple reactants (elements or products) form a single compound.
General form: A+B AB

Ex)

Decomposition (3): be able to identify the type; no prediction

A reaction where one complex compound breaks into multiple simple products
Opposite of synthesis
General form: AB A+B

Ex)

Single Replacement/ displacement (4): be able to do simple product prediction (not if it occurs)

General form: A + BC  AC + B

Ex)

TIP: Cations switch with other cations, and anions switch with other anions.
Driven by the reactivity of elements

Double Replacement/ displacement (5): be able to do complex product prediction and if the reaction occurs.

A reaction where two ionic compounds switch their cations and anions

General form: AB +CD  AD +CB

Ex) precipitate:

Ex) Acid-base:

Driven by the formation of a solid, liquid, or gas

Usually at least one of your reactants is a solution (aq: dissolved in water)

This allows compounds to ionize (separate into ions) and then switch their ions

Predicting products

Double Replacement

1)Predict products

2)Check subscripts

3)Label States {TIP: Soluble (aq), insoluble (s)}

4)Balance (TIP: if all products are aq, there is no reaction)

Combustion

is always the reactant
____ + +

Redox reaction

Oxidation-reduction (redox) have two processes that take place:

Oxidation: The loss of electrons. ex- Rust
Reduction: The gain of electrons. Ex- silver mirror, electroplating

These two processes always take place together

OIL RIG: Oxidation Is Lost, Reduction Is Gained

LEO the lion says GER: Lose Electrons is Oxidation, Gain Electrons is Reduction

Any reaction where electrons are transferred is a redox reaction

How so we tell when redox occurs? Oxidation number/state

The oxidation number of an atom/ion helps us keep track of electrons in a reaction.
Atoms/ions are oxidized or reduced, not compounds (An element not a compound)
Atoms/ions that become more positive are oxidized (Cu Cu2+)
Atoms/ions that become more negative are reduced (Cu2+  Cu)

If oxidation numbers don’t change, there is no redox reaction.

Oxidation Number/State Rules:

1)The oxidation number for an atom in its elemental form is always zero.

A substance is element if both of the following are true:
Only one kind of atom is present
There is no charge on the element
Subscripts don’t matter, the elements just have to be by themselves.

2)The oxidation number of a monoatomic (single type of atom) ion= charge of the monatomic ion.

3)The oxidation number of all group 1 metals is 1+ (unless elemental).

4)The oxidation number of all group 2 metals is 2+ (unless elemental).

5)Hydrogen (H) has two possible oxidation numbers:

1+ when bonded to a nonmetal (covalent compounds)
1- when bonded to a metal (ionic compounds)

6)Oxygen (O) has two possible oxidation numbers:

2- in almost all compounds (peroxide is the exception)
1- in peroxides… very rare don’t need to worry about it

7)The oxidation number of fluorine (F) is always 1-.

8)The sum of the oxidation numbers of all atoms (or ions) in a neutral compound is 0.

9)The sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge on the polyatomic ion.

Balancing

Redox reactions must be balanced by both mass and charge because electrons are transferred during the reaction.
Balancing by charge conserves the amount of electrons transferred in the reacting
To balance a redox reaction, we use the half reaction method
Half Reaction Method: breaks the whole reaction into two parts

1)Oxidation half reaction: electrons go on the right (product)

2)Reduction half reaction: electrons go on the left (reactant)

Half Reaction Method Steps:

1)Write separate equations for the oxidation and reduction half- reactions.

2)For each half reaction:

Balance all the elements except hydrogen and oxygen
Balance oxygen using
Balance hydrogen using H+
Balance charge using electrons

3)If necessary multiply one or both balanced half-reactions by an integer to equalize the number of electrons transferred in the two half-reactions.

4)Add the half- reactions, and cancel identical species.

5)Check that the elements and charges are balanced.

Conversions and Stoichiometry

Look at the handout from teacher!!!

Solutions and Molarity

Solutions:

A solution is a mixture of two or more substances that are homogenous (the same throughout)
Solvent: the substance you have more of. The thing the substances are dissolved in. Usually water.
Solute: The substance you have less of. The substance that is dissolved.
All solutions are made up of a solvent and a solute.
Solvents and solutes can be liquids, solids, or gasses.

Ex) Salt water: Solvent- H2O (water) Solute- salt

Ex) Soda: Solvent- H2O (water) Solute: sugar, dye, CO2 (carbon dioxide)

Concentration: amount of solute dissolved in an amount of solvent

Can be measured many different ways (ppm, Molarity) but the most common is molarity (M)

Moles: of solute Liters: of solution

Calculating Molarity:
Ex) 0.10 mol. of NaCl is dissolved in 1.0 L of H20. What is the molarity of the solution?

Ex) What is the concentration of a solution where 16.0g of sugar, , is dissolved in 250 mL of water?

TIP (for above problem): Always divide by volume.

Saturated Solution: Maximum amount of solute possible has dissolved in a solvent.

Unsaturated Solution: Less than the maximum has been dissolved.

Supersaturated Solution: More than twice the maximum has been dissolved.

A more concentrated solution means that there are more molecules per unit volume. Two solutions with the same concentration, but with different volumes have the same number of molecules per unit volume, but a different total number of molecules.

Preparing Solutions

Two Methods for making a solution:

1)From a solid ( dissolve solid in H2O)

2)Diluting an already made solution

Preparing from a solid (1)

Ex) you want to make a 0.15 M NaCl solution and you need 2.0L of that solution. What is the mass of NaCl you will need?

** M*V(in liters)= Mol.  g.