final exam study notes

Final Exam Study Notes

CP ChemistryPeriod 5

Physical and Chemical Properties

Physical Properties Chemical Properties• Size• Color• Texture• Smell• Mass• Taste• Density• Volume• Area• Melting Point• Malleability• Elasticity• Solubility

• Flammability• Digestibility• Decomposable• Ability to oxidize• Reactivity inert (not reactive)

Physical Property- a trait or characteristic that you can observe without changing the identity of the substanceChemical Property - a trait you can observe by changing the identity of the substance

Physical and Chemical Changes

Physical Changes Chemical Changes

Examples in terms of a piece of paper• Crumpled paper• Ripped paper• Drawn on paper• Stomping on paper

Examples in terms of a piece of paper• Eating it• Digesting it• Burning it

Physical Change- a change that affects only physical properties and does not alter the identity of the substanceChemical Change- a change that alters the identity of our substance

Elements, Compounds, and Mixtures

• Element- a substance that cannot be separated by chemical or physical means

• Compound- a substance made up of two or more elements only separated by chemical means

• Atom- smallest unit of an element• Molecule- smallest unit of a compound

Elements, Compounds and Mixtures cont.

• Mixture- a combination of substances thhat are not chemically combined. These can be separated physically– Homogeneous: looks same throughout– Heterogeneous: composed of different parts

Homogeneous Heterogeneous

Accuracy and Precision

Trial 1

Trial 2

Trial 3

Trial 4

Trial 5

0

0.5

1

1.5

2

2.5

3

Mass 1Mass 2

True Data is accurate.Repeatable data is precise.

Accurate & Precise

Counting Protons/Neutrons/Electrons

The mass number of the element eqauls the number of protons in an element

The number of protons is also the number of electrons unless its an ion

To find the neutrons you subtract the atomic number from the mass number

Isotopes Are atoms that have lost or gained

neutrons, same element but different number of neutrons

Radiation particles/ Nuclear equations

Alpha- 42H- stopped by clothing or skin

Beta- 0 -1 e- stopped by a sheet of lead

Gamma- stopped by several inches of lead, most dangerous

Nuclear reactions happened when there is an unstable particle and eventually gives off a particle of radiation

Ions Atoms that have lost or gained electrons

atoms turn into ions when electrons move Ions have a charge

There are negative electrons and positive electrons

How Ions are formed: Positive ions have lost electrons Negative ions have gained electrons Positive ions are called cations Negative ions are called anions When atoms are most stable they have an octet

Octet- 8 electrons in the outer most energy level

Covalent Bonds Share electrons between two atoms Properties

Low melting point Molecule structure Gases, liquids, soft solids Poor conductors of heat Poor conductors of electricity Typically 2 non-metal atoms

Ionic Bonds They trade electrons between the two

atoms Ions must form from the atom

Properties High melting point Crystal lattice structure Hard solids Brittle Good conductors of heat Good conductors of electricity Typically 1 metal and 1 non metal

Wavelength

• Waves of Light: electromagnetic radiation (light) moves as a wave

Crest

Trough

Wave

Wavelength/Frequency

• λ= Wavelength: Distance from crest to crest on a wave

• v= Frequency: How often a wave passes by in a second (s-1)

• Wavelength and Frequency are inversely related – Wavelength increases, Frequency decreases – Wavelength decreases, Frequency increases

Calculations

• E=h• E= Energy (Joules)• h= Plank’s Constant (6.626 x10-34) • = Frequency

• Example: A yellow light has a wavelength of 600nm – a) What is the frequency of the light?– b) What is the Energy of the light? • Answers on next slide

Answer – Frequency

• a) = 600nmn v= x c= 3.0 x 108 m/s

= c/v 600nm x 1n/10 x 9nm= 6.0 x 10-7m

6.0 x 10-7m= 3.0 x 108 m/s /vV= 3.0 x 108 m/s / 6.0 x 10-7m = 5.0 x 1014 s-1

Answer- Energy

• E=hv • E= ?• h= 6.626 x10-34 J(s) • v= 5.0 x 1014 s-1 • E= (6.626 x10-34 ) 5.0 x 1014

– 3.31 x 10-19 J

Ionization Energy• The amount of energy

needed to remove one electron from an atom

• Size of atom determines how easily electrons are removed

• Big atoms lose e- with minimal effort

• Little atoms lose e- with a huge amount of energy needed

- Increases up and to the right on the Periodic Table

• Noble gases all have elect negativities equal to zero

• If an atom needs a lot of energy to remove an electron its because it really wants the e-

Periodic Trends

• Properties of elements can be predicted using the location on the periodic table – Electron Configuration – Family and Periods – Densities – Reactivity – Atomic radius – Ionization Energy – Electronegativity

Atomic Radius

• The distance from the nucleus to the outer-most elections (in the highest energy orbital filled)

• As electrons fill into higher energy orbitals, the radius of the atoms gets bigger!

NaK

Rb

Na 3s1

K4s1

Rb5s1

Atomic Radius

• Within an energy level adding more protons makes the radius of atoms smaller because the protons can hold the electrons in closer

PS

Cl

P3P3 S

3 P4Cl

3P5

Question

• Put in order smallest to largest: – Rb, P, Na • Answer on next slide

Answer

• Na, Rb, P

Periodic Table Families and Periods• Groups (families) = The columns on the Periodic Table• Periods = Rows on the Periodic Table • Elements arranged by atomic number

– Column 1: #1 • Silvery White

– Column 2: • React with water to form a base

– Column 3: 5,8,9,10,11,12 • All metals • Form colorful solutions • Hard Brittle • Metallic • Versatile in bonding ability • Charges: +2, +3

Periodic Table Families and Periods• Column 14

– Charge +4 -4 – Non-metallic – Can bond with positive or negative

ions – Solid at room temperature – Relatively low reactivity

• Column 16: Oxygen Family– Charge: -2 – Non-metallic – Bonds with itself– Shares electrons with other

elements

• Column 17: Halogens- Non-metallic - Charge -1- Very reactive with positive ions - Not solid at room temperature - Colored Gas

• Column 18: Noble Gases - Non-Metallic - Non Reactive - Charge: 0

VSEPR Molecule Geometries

VSEPR is a model of molecular structures based on the idea that ideal structures minimize electron pair repulsions

Used to draw and evaluate Lewis Structures

Bare electrons are the most repulsive!

Molecular Geometry Models Looking at the molecular geometry

of a single atom, not of an entire molecule

3D Figures to represent Lewis Structures

Constituent groups are the things bonded to the atom under scrutiny

Dashed lines represent a bond behind the plane of the paper; wedged lines represent a bond coming toward you

Planar Geometry

• Linear• 1-2

Constituents• 0 Lone Pairs• Bond Angle:

180

• Trigonal Planar

• 3 Constituents

• 0 Lone Pair• Bond Angle:

120

• Bent• 2 Constituents• 1 Lone Pair• Bond Angle:

<120

Tetrahedral & Derivatives

• Tetrahedral• 4 Constituents• 0 Lone Pair• Bond Angle:

109.5

• Trigonal Pyramidal

• 3 Constituents• 1 Lone Pair• Bond Angle:

107.3

• Bent• 2 Constituents• 2 Lone Pairs• Bond Angle:

104.5

Lewis Dot Diagrams

Lewis dot structure is a drawing of how the atoms are bonded together covalently using valence electrons.

You need to know Shared pair= 2 electrons shared by 2

atoms (bond) Lone pair= 2 electrons not shared by

atoms (unshared pair)

How to draw Lewis Dot Structure

1. Count the total valence electrons for the moleculeEx: SCl2=20 valence electrons

2. Select a central atom.look for= *the only one of its kind.

*less electronegativeEx: SCl2= S is central atom because it’s alone

How to draw Lewis Dot Structure

3. Set up the elements as symmetrical as possible. Ex: SCl2= Cl S Cl

4. Draw in shared pairs by drawing a line. Ex: SCl2= Cl-S-Cl

How to draw Lewis Dot Structure

5. Account for electrons used from total you started with.(Shared pairs=2 electrons) Ex: 20 valence electrons -4 shared electrons ___ 16 unshared electrons

How to draw Lewis Dot Structure

6. Fill in unshared pairs around the outside of elements Ex:

How to draw Lewis Dot Structure

7. When there isn’t enough electrons for every element to have an octet, we share more pairs.

Ex:

How to find polarity in molecules When finding the polarity in

molecules you need to find out if the bonds are polar or non-polar first. Polar bond- when 2 atoms share

electrons unequally Non-Polar bond- share electrons

completly even.

How to find polarity in molecules Polar Molecules- Molecules where

one side of the molecule has more electrons than the other. 1. if there is a lone pair on the center

atom, it is polar 2. If bonds are unequal polarity, than the

molecule is polar

How to find polarity in molecules Non-Polar Molecules

If there is no lone pairs on the center atom

If the bonds are equal polarity

5 Types of Chemical Reactions • Synthesis:• Ex: Cu+3 + O2 Cu2O3

• Decomposition• Ex: MgCl2 Mg+2 + Cl

• Single Replacement• Ex: MgCl2 + Cu+2 Mg + CuCl2

• Double Replacement• Ex: 3MgCl2 + Cu2O3 3MgO + 2CuCl3

• Combustion• Hydrocarbon + oxygen CO2 +H2O• CH3OH + O2 CO2 + H2O

Predicting Products (using the 5 type of chemical reaction)• Synthesis all you have to do is combine the two reactants

to get your products• Decomposition you break up your reactants and get two

products• Single replacement take either the positive or negative

ion (by itself) and replace it with the positive or negative ion from a formula in the equation

• Double replacement you take the positive ion from one formula and put the negative ion from the other formula to create a new formula, do this again with the left over positive and negative ions. (positive come first)

• Combustion always ends up with carbon dioxide and water

Law of Conservation of Matter• Matter can neither be created nor destroyed

• In chemical equation it’s crucial to make sure its balanced because, if its not balances it goes against the law of conservation of matter because it creates (or destroys) matter

Balancing Equations• According to the law of conservation of matter you

have to balance all of your equations so that you don’t create or destroy matter.

• NaOH + Cl2 NaCl + OH• This is not balance because you have two

chlorines in the reactants and only one on the product side

• So, all you have to do is add a coefficient in order to balance it:

• 2NaOH + Cl2 2NaCl + 2OH

Net Ionic Equations• Aqueous substance dissociate• A complete ionic equation shows all the ions and

molecules in a reaction• Zn (s) + CuSO4 (aq) ZnSO4 (aq)• Complete ionic equation:• Zn (s) + Cu+2 (aq) + SO4-2 (aq) Zn+2 (aq) + SO4-2 (aq)

+Cu (s)• NET IONIC:• Anything that’s AQUEOUS that’s the same on both sides

you can cancel out• So you can get rid of SO4-2 (reactant side) AND SO4-2

(product side)• Final net ionic equation:• Zn (s) + Cu (aq) Zn (aq) + Cu (s)

Problems• 1. What type of reaction is this?

• Mg + KCl MgCl + K

• 2. Balance the following equations:• ___ Al + ___ O2 _____ Al2O3 • ___CuS + ____ O2 ____CuO + _____ SO2

• ____ Ca3P2 + ____ H2O ____ Ca(OH)2 + ___ PH3

• 3. Write the net ionic equation for:• AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq)

ANSWERS• 1. single replacement

• 2. 4 Al + 3 O2 2 Al2O3

• 3. 2 CuS + 3 O2 2 CuO + 2 SO2

• 4. (1) Ca3P2 + 6 H2O 3 Ca(OH)2 + 2 PH3

• 5. AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq)

• Complete ionic = • Ag+1 (aq) + NO3

-1 (aq) + Na+1(aq) + Cl-1(aq) Ag+1 (aq) + Cl-1 (s) + Na+1 (aq) + NO3

-1 (aq) • Net Ionic =• Ag+1 (aq) +Cl-1 (aq) AgCl (s)

The Factor Label Method

• A ratio used to convert the unit you have into the desired unit

• Example: If you are given one day, how do you convert it into the amount of seconds in a day?

Answer:1 day* 24 hours*60 minutes*60 sec

1 day 1 hour 1 min

Cross cancel the units!!

The Factor Label Method Cont.

• The factor label method is useful in converting metric prefixes

• The Metric Prefixes are: – TGMKHDBDCMMNP– The Great Mister King Henry Died By Drinking Chocolate Milk Monday Night Partying

– Tetra Giga Mega Kilo Hecto Deca BASE Deci Centi Mili Micro Nano Pico

You can use Metric conversions to change from prefix to prefix!

Converting Moles, Liters, Grams & Particles

• 1 Mole = 6.02 X 1023 “things”– 6.02 X 1023 = Avogadro’s Number

• Moles to Particles/Liters/Grams: mole of element X 6.02*1023 = # with desired unit

1 Mole

Molar Mass

• To find the molar mass you must refer to the periodic table

• Look up each atomic mass of the element and add them all together to find the molar mass

• If there is a subscript then you must multiply the atomic mass by the subscript

Empirical Formula

• Empirical formula= the lowest terms ratio of elements in a formula (Not the true formula)

• To calculate the empirical formula you must find 1. Percent to Mass2. Mass to Mole3. Divide by small4. Multiply until whole

Molecular Formula

• Molecular Formula= The true formulas for a compound (Not lowest terms)

• To find the molecular formula you must divide the actual molecule mass by the empirical formula mass– Example: OH is the empirical formula, the actual formula has a mass

of 34 g/mol, what is the molecular formula?– 34 g/mol = 2 so OH becomes O2H2

17 g/mol

Percent Composition

• To calculate the % composition you must use the following equation:

Mass of particle * 100 Mass of whole

States of Matter

Gas Forms to shape of container Spread apart atoms

Solid Atoms tightly compacted Definite shape and volume

Liquid Definite volume Moderate spacing of atoms

Parts of SolutionsSolute

Dissolved by the solvent in the solution Ex: Salt in salt water

Solvent Substance that does the dissolving Ex: Water in salt water

Phase Diagrams/ Heating Curves

Phase Diagrams Shows what temperature and pressure combinations

can create each state of matter for a particular chemical.

Heating Curves Shows the temperatures at which changes in states of

matter occur and describe how a substance uses heat.

Molarity Calculations

Molarity= moles of substance/ volume(L) Ex: The chemical Carbon Dioxide has a volume of 2L. Find

the concentration of Carbon Dioxide if it has a mass of 24.02g

Grams to Moles using Molar Mass conversions. 24.02g*1mol/12.01g/mol= 2mol

Then use the formula M(Molarity)=mol(Moles)/Volume(L) M=2mol/2L M=1M

Intermolecular Forces

Intermolecular Forces The forces of attraction between molecules.

Vander Waal's(London Dispersion)Hydrogen BondingDipole-Dipole

SPECIFIC HEAT

Specific Heat Capacity- The amount of heat needed to raise the temperature of one gram of substance one degree Celsius or one degree Kelvin.

Molar Heat Capacity- The amount of heat needed to raise the temperature of one mole of substance one degree Celsius or one degree Kelvin.

SPECIFIC HEAT CALCULATIONS

q = m C THeat (Joules)

Mass (g)or moles (mols)

Specific or molar heat capacity

Change in temperature

(kelvin or celsius)

Tf – Ti

EXAMPLE PROBLEM

12g of water is heated from 15 degrees Celsius to 35 degrees Celsius. How much heat was absorbed by the water? (The specific heat capacity for water is 4.184 J/g degrees Celsius).

Q = ?

m = 12 g

C = 4.184 J/g

T = 20 degrees Celsius

q = 12 * 4.184 * 20

q = 1004.16 J

PROPERTIES OF ACIDS

Sour

Burn/ sting

React with metal

Electrolyte (conducts electricity)

pH less than 7

Releases hydrogen ions in water

Accept an electron pair

PROPERTIES OF BASES

Bitter

Slippery

Non-reactive with metals

Electrolyte

Releases hydroxide ions in water

Donate an electron pair

PH CALCULATIONS

The pH scale is a numerical system that expresses the acidity of a solution

What is the pH of a solution that has [H+]

of 1.38 * 10-11 ?

(plug into your calculator)

Ans: pH = 10.86

POH CALCULATIONS

A numerical scale that measures solutions by basicity.

pOH = -log [OH-]

What is the pOH of the solution you used in the last slide?

(plug into your calculator)

Ans: 3.14

Hint: pH + pOH = 14

H+ CALCULATIONS

[H+] = 10-pH

Example:

Determine the concentration of [H+] in the solution. pH = 3.0

Plug into your into your calculator by clicking “2nd” and log

Ans: 1 x 10-3 M

[OH-] CALCULATIONS

[OH-] = 10-pOH

Example:

Determine the [OH-] in the solution given the pOH is 4.0.

Ans: 1 x 10-4 M

Kinetic Molecular Theory Definition: A theory concerning the

thermodynamic behavior of matter, especially the relationships among pressure, volume, and temperature in gases.

Absolute Zero and STP STP is used for measuring gas

temperature and volume. STP means standard temperature pressure.

Absolute zero is used for Absolute zero is the point where no more heat can be removed from a system, according to the absolute or thermodynamic temperature scale. This corresponds to 0 K or -273.15°C.

Measuring pressure Pressure is

measured using a barometer. The glass tube on the barometer contains a vacuum that allows mercury flow up it when pressure is excreted on the surface of the mercury

Pressure is usually measured in:

Atmospheres (atm) Bar (bar) Pascals (pa) Millimeter of

Mercury (mmHg) Torr (torr)

Pressure ConversionsPRESSURE CONVERSION PROBLEM

1 atm= 101.325 Pa 1 bar= 100.025 Pa 1 Torr= 133.32 Pa 1 MMHg= 133.32 Pa 1 MMhg= 1 Torr

A radio station announcer reports the atmospheric pressure to be 99.6 kPa. What is the pressure in atmospheres? In millimeters of mercury?

Pressure Problem answer99.6 kPa x 1 atm/101.3 kPa = 0.983 atm0.983 atm x 760 mm Hg/1 atm = 747 mm HgAnswer0.983 atm; 747 mm Hg

Gas Laws Gas laws describes observed behaviors of

gasses.

Charles law equation: Gay-lussac law equation:

More Gas Laws Grahams Law: Ideal Law: Dalton Law:

Gas Law Problems 1) In a thermonuclear device, the pressure of 0.050 liters of gas within the bomb casing reaches 4.0 x 106 atm. When the bomb casing is destroyed by the explosion, the gas is released into the atmosphere where it reaches a pressure of 1.00 atm. What is the volume of the gas after the explosion?

2) On hot days, you may have noticed that potato chip bags seem to “inflate”, even though they have not been opened. If I have a 250 mL bag at a temperature of 19 0C, and I leave it in my car which has a temperature of 600 C, what will the new volume of the bag be? 

Gas Law Answers 1) 2.0 x 10 5 L

2) 285 mL