Notes Chemistry Pre-Comp Pt. 1 2013 Significant ... - Weebly

Notes Chemistry Pre-Comp Pt. 1 2013 **For blank sections you can either ask me or use your notes!!

1 Made By Ashley Thomas

i. Significant Figures

Types of Zeros

o Sandwich zero: 505

o Leading zero: 0.0005

o Trailing zero: 5000

Rules

o All non-zero digits are significant

o All sandwich zeros are significant

o Leading zeros are never significant

o 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.

ii. Data Analysis

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)

Know the dartboards!!!!

iii. Physical and Chemical Changes

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.

iv. The Periodic Table

Metals vs. Metalloids vs. Non-metals

o Metals: to the left of the stairstep

o Non-metals: to the right of the stairstep

o Metalloids: Directly touching the stairstep



Groups

o 1a/b-Alkali Metals

o 2a/b- Alkaline Earth Metals

o 3b-12b- Transition Metals

o 7a/17b- The Halogen

o 8a/18b- Noble Gasses

o FOR CHARGES SEE PERIODIC TABLE 1

Rows are called periods (n).

Periods are the highest energy level of an elements electron configuration.

The atomic number tells you the total number of electrons.

3- energy level p- subshell - number of electrons

v. Atomic Structure

Protons: located in nucleus, positive charge (+), equal to atomic number

Neutrons: located in nucleus, no charge, equal to atomic mass minus atomic number

Electrons: located in orbital’s (around nucleus), negative charge (-), in a neutral atom equal to

number of protons

Electrons

o We know where electrons are based on their electron configuration.

o Quantum Numbers

n: Principle quantum number/ energy level

: (s,p,d,f) shells/sublevels

: orbitals

: Electron Spin

o Each sublevel is made up of orbitals, each orbital is two electrons.

S: 1 orbital

P: 3 orbitals

D: 5 orbitals

F: 7 orbitals

Orbital diagrams: visual representation of the quantum number.

Steps to draw a orbital diagram:

1) Write the electron configuration (longhand).

2) Draw boxes/lines representing each orbital of the electron configuration.

3) Label your subshells.

4) Fill electrons according to rules listed below.

Pauli Exclusion Principle: No electron may have the same four quantum numbers.

Aufbow Principle: Fill orbitals from lowest to highest energy level.

Hund’s rule: Every orbital must be singly occupied by one electron before any orbital is doubly

occupied. All singly occupied orbitals have electrons with the same spin. Reason: Electrons repel

each other. This only applies to p, d, and f subshells.

vi. Electron Configuration



Sub-shells (s,p,d,f)

o S: Groups 1-2, max electrons= 2

o P: Groups 13-18, max electrons= 6

o D: Groups 3-12, max electrons= 10

o F: Groups Lanthanum/ Actinium Series, max electrons= 14

Noble Gas/ Short hand configuration steps:

1) Find the element on the periodic table.

2) Write the noble gas from the previous period in brackets.

3) Start writing the normal electron configuration from the start of the period the element is

in.

vii. Periodic Trends

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.

o Shielding electrons “shield” the outermost electrons from the nucleus “pull”.

o Called effective nuclear charge ( )

Z is the atomic number or number of protons

S is the number of shielding electrons

o The higher the the more the valence electrons are attracted to the nucleus.

o The higher the the more likely the element will become an anion(negative ion).

Periodic Trends (See Periodic Table 2)

o 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.

o Atomic Radius/size: see above notes

o 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”.

o Electronegativity/Electron Affinity: How much an atom attracts an electron.

o 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.

o Metallicity/Metallic Character: How much like a metal an element is.

o 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:

o : effective nuclear charge (only explains trends across a period); every atom in the same

group has the same

o Distance/Size: The farther from the nucleus, the less “pull” the nucleus can exert. (Columb’s

Potential Energy)

o 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.

viii. Nomenclature: Naming compounds

Charges:

o Zn (zinc) is always 2+

o 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.

o 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.

o Binary: Name of metal + name of non-metal plus –ide.

o Polyatomic: name of metal + mane of polyatomic ion.

o 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.

o HCl: Chloride hydro-chlor-ic acid

Hydro-name of anion- ic +acid

TIP: Hydro- only H and 1 element

o : Chlorite Chlor-ous acid

Name of anion- + acid

o :Chlorate Chlor-ic acid

Name of anion-ic + acid

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

ix. 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

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

o 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

o 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

x. Percent Composition- by mass

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

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

xi. Molar Mass:

The mass of one mole of a substance

Steps:

1) Find the mass of each element (found on the periodic Table)

2) Multiply each element by its subscript.

3) Add all the masses together

TIP: DON’T FORGET TO DISTRIBUTE!!!

xii. Empirical and Molecular 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



.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

o How to convert between empirical and molecular formulas:

Multiply subscripts of empirical by the whole number

o 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?

xiii. Chemical Equations

A chemical equation is the “recipe” for a chemical reaction

Parts of an Equation:

o Reactants: left side of the arrow, “ingredients”



o Products: to the right of the arrow, what you make

o Coefficients: The amounts (units=moles) of each compound.

o Subscripts: the amount of each type of atom/element

o States: (s) (l) (g) (aq)

(s): solid

(l): liquid

(g): gas

(aq): aqueous solution (in water, compound exists as separate ions) (in Latin:

surrounded by water)

Examples:

1)

2 moles of solid potassium reacts with 1 mole of chlorine gas to form/produce 2 moles of

solid potassium chloride.

2)

1 mole of solid zinc reacts with 1 mole of an aqueous solution of lead (II) nitrate to form

1mole of an aqueous solution of zinc nitrate and one mole of solid lead.

3) 1 mole of a calcium hydrogen solution reacts with 2 moles of a hydrochloric acid solution to

produce 1 mole of calcium chloride solution and 2 moles of liquid water.

Diatomic elements: H, N, F, O, I, Cl, and Br all exist as diatomic molecules, not individual

elements.

o , , , , , ,

o Memory Devise:

HOFBrINCl twins

xiv. Balancing Reactions

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?

o Change the coefficients of compounds

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

o Change the subscripts of an element

o Change the elements

TIPS:

o Treat polyatomic ions as a unit or group

o Write water (H2O) as HOH

o No fractions in final answers

o Reduce coefficients only if you can reduce all of them

o For combustion reactions balance in the following order: H C O

xv. Types of Reactions

5 basic types:

1) Combustion



2) Synthesis

Oxidation- reduction

3) Decomposition


4) Single Replacement


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.

o 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

o A reaction where simple reactants (elements or products) form a single compound.

o General form: A+B AB

Ex)

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

o A reaction where one complex compound breaks into multiple simple products

o Opposite of synthesis

o General form: AB A+B

Ex)

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

o General form: A + BC AC + B

Ex)

Ex)

o TIP: Cations switch with other cations, and anions switch with other anions.

o Driven by the reactivity of elements

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

reaction occurs.

o A reaction where two ionic compounds switch their cations and anions

o General form: AB +CD AD +CB

Ex) precipitate:

Ex) Acid-base:

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

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

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

o 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)

o 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?

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

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

o Unsaturated Solution: Less than the maximum has been dissolved.

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

o 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.

o 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.



xvi. Redox Reactions

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)

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

o 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.

xvii. Balancing Redox Reactions

o Redox reactions must be balanced by both mass and charge because electrons are transferred

during the reaction.

o Balancing by charge conserves the amount of electrons transferred in the reacting

o To balance a redox reaction, we use the half reaction method

o 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)

o Half Reaction Method Steps:



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

2) For each half reaction:

a. Balance all the elements except hydrogen and oxygen

b. Balance oxygen using

c. Balance hydrogen using H+

d. 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.

xviii. Predicting Products

o 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)

o Combustion

is always the reactant

____ + +

xix. Solubility Rules

Ions that Form Soluble Compounds Exceptions

Group 1 ions (Li+, Na+, ect.) N/A

Ammonium N/A

Nitrate N/A

Acetate N/A

Bicarbonate (aka Hydrogen Carbonate) N/A

Chlorate N/A

Perchlorate N/A

Halides (Cl-, Br-, I-) When combined with Ag+, , and

Sulfate When combined with Ag+,

Ions that Form Insoluble Compounds Exceptions

Carbonate When combined with always soluble ions

Chromate When combined with always soluble ions, , or

Phosphate When combined with always soluble ions



Sulfide When combined with always soluble ions

Hydroxide When combined with always soluble ions,

xx. Polyatomic Ions

Nitrate

Chlorite

Nitrite Hydrochlorite

Chromate

Periodate

Dichromate

Iodate

Cyanide Hypoiodate

Hydroxide Bromate

peroxide Hyprbromite

carbonate

Bicarbonate*

Sulfate

Bisulfate**

Sulfite

Permanganate

oxalate Hydrogen Sulfide

Phosphate

Borate

Phosphite

Silicate

Thiosulfate

*** Acetate

Perchlorate

Ammonium

Chlorate

*AKA Hydrogen Carbonate

**AKA Hydrogen Sulfate

***Can also be written as

Notes Chemistry Pre-Comp Pt. 1 2013 Significant ... - Weebly

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