Chemistry Honors Semester 1 Study Guidechemhonorssemester1.wikispaces.com/file/view... · Chemistry...

Good luck studying guys. –Sandeep :)

Chemistry Honors

Semester 1 Study Guide

Chemistry Honors Study Guide| Notes 2

Introduction to Chemistry: Ch1

Key Terms:

Chemistry the study of matter and the changes it undergoes

Substance matter that has a definite and uniform composition, (also called a chemical)

Model visual/verbal/mathematical explanation of experimental data

Mass measurement that reflects amount of matter

Weight measure of matter and the effect of Earth’s gravitational pull on that matter

Conclusion judgment based on information obtained

Control a standard for comparison

Dependent Variable value that changes in response to change in independent variable (often y-axis)

Experiment set of controlled observations that test the hypothesis

Hypothesis tentative explanation for what has been observed

Independent Variable value that you intend to change (often x-axis)

Qualitative Data information that describes a physical characteristic; relating to 5 senses

Quantitative Data numerical description; tells how much, how long, how fast, etc.

Scientific Law relationship in nature that is supported by many experiments

Scientific Method organized approach to doing an experiment (see 1.3, Scientific method)

Applied Research research undertaken to solve a specific problem

Pure Research done to gain knowledge for the sake on knowledge itself

Theory explanation of a natural phenomenon based on many observations and investigations over time

Notes

Ozone Layer

When O2, oxygen gas, is exposed to ultraviolet radiation, O3, ozone, is formed. In the past, Chlorofluorocarbons (CFC’s)

were used in refrigerators and air conditioners, as well as plastic foams and propellants in aerosol cans. Scientists

discovered that CFC’s had gone into the atmosphere and bonded

with the O2, preventing it from being formed into ozone and leading

to the lower levels of ozone we have in our atmosphere today.

1.3, Scientific Method

The scientific method is a systematic approach to experimentation.

First, make an observation: when ever I throw something, it comes

down. Next, ask a question: if I throw this ball up, it will come down?

Then, make a hypothesis, which would be what you think will

happen: if I throw this ball up, it will come down. Predict what will

happen: a ball will fall back down. Test experiment: it worked!

Finally, make more predictions and test them out: A rock will also

fall. However, if you were wrong, go back and revise your

hypothesis, then test out the experiment again.


Analyzing Data: Ch2

Key Terms:

Base unit defined unit based on object in physical world: second, meter, kilogram, mol, etc.

Density amount of mass per volume (formula below)

Derived Unit unit that is a combination of base units: volume, density, speed, etc.

Kilogram SI unit for mass

Kelvin SI unit for temperature, at 0 Kelvin, molecular motion stops (formula below)

Liter measure of volume, 1 L= 1 dm3

Meter SI unit for length

Second SI unit for time

Conversion Factor ratio of equivalent values having different units.

Dimensional Analysis systematic approach to problem solving that uses conversion factors.

Scientific Notation used to express any number as a number between 1-10 times 10 to a power.

Accuracy how close a measurement is to an accepted value.

Precision how close a series of measurements are to each other

Error difference between experimental value and accepted value

Significant Figure include all known digits plus 1 more. (explained below)

Graph visual display of data

Formulas:

Density =

Error = experimental value – accepted value

Error

Percent Error = ----------------------------

Accepted value

y2 - y1

Slope = ----------------------------

x2 - x1

Mass ----------------

Volume


Analyzing Data: Ch2 –cont.

Notes The SI (Systeme Internationale d’Unites) is a French system that updated the metric system. It is used between

scientists of different nations in most of the world.

A good guide to dimensional analysis can be found at

http://www4.ncsu.edu/unity/lockers/users/f/felder/public/kenny/papers/units.html.

Accuracy vs. Precision

Graphs

Pie charts show parts of a whole. Bar graphs show how a factor varies with time, location, or temperature.

Information about significant digits can be

found at the very end, in the chemistry

skills section.

http://www4.ncsu.edu/unity/lockers/users/f/felder/public/kenny/papers/units.html


Matter—Properties and Change: Ch3

Key Terms:

states of matter different physical forms of all matter that exists on earth

solid definite shape and definite volume

liquid flows, has constant volume, take shape of container

gas flows to conform to shape of container and fills entire volume

vapor gaseous state of a substance that is sold or liquid at room temperature

physical property prop. that can be observed/ measured without changing the sample’s composition

chemical property the ability of a substance to combine with or change into one or more other s

extensive property dependent on the amount of a substance; length, volume, mass

intensive property independent of the amount; density, scent

physical change alters a substance without changing its composition; cutting, crushing

chemical change process that involves one or more substance changing in to new ones; chemical reactions

phase change transition of mass from one state to another

law of conservation of mass mass is neither created nor destroyed during a chemical reaction

mixture combination of two or more pure substance, both retain individual chem. properties

heterogeneous mixture mixture in which individual substances remain distinct; salad dressing, pulpy juice

homogeneous mixture mixture that has constant composition throughout

solution same as homogenous mixture

filtration uses porous barriers to spate solids from liquids

distillation uses different boiling points of substances to separate them

crystallization results in the formation of solid crystals from a mixture; rock candy

sublimation separating by changing solid substances directly into vapors

chromatography separates components of mixture based on the ability of each to travel across a surface(marker lab)

element pure substance that cannot be separated into simpler substances by physical or chemical means

periodic table ment organizes elements into grid of horizontal rows (periods) and vertical columns (families)

compound two or more different elements that are combined chemically

law of definite proportions compounds are always composed of the same elements in the same proportion by mass

law of multiple proportions when two compounds are made of the same elements, they will have whole number ratios

percent by mass ratio of the mass of each element to the total mass of the compound expressed as a

percentage

Formulas:

Mass element

Percent by Mass = ---------------------------- X 100 Mass compound


Matter—Properties and Change: Ch3 –cont.

Notes

Law of Definite Proportions

The law stating that a pure substance, e.g. H2O, will

always have the same percent by weight, e.g. 11.2%

H and 88.8% O. In other words, oxygen will always

make up about 88% of any amount of water.

Law of Multiple Proportions

When two or more elements form more than one

compound, the ratio of the weights of one element

that combine with a given weight of another

element in the different compounds is a ratio of

small whole numbers. For example, carbon and

oxygen combine in carbon dioxide (CO2) and carbon

monoxide (CO). A sample of carbon dioxide

containing 1 gram of carbon contains 2.66 grams of oxygen; a sample of carbon monoxide containing 1 gram of carbon

contains 1.33 grams of oxygen. The ratio of the two weights of oxygen (2.66:1.33) is exactly 2:1. So therefore there are

twice as many oxygen atoms in a molecule of carbon dioxide compared to carbon monoxide.

The law of Multiple Proportions is definitely one that you would want to do practice problems in the book for.


The Structure of the Atom: Ch4

Key Terms:

Atom smallest particle of an element that still retains its properties

Cathode Ray Radiation that travels through a cathode ray; these were used in TVs.

Electron negatively charged particles that travel around the nucleus

Nucleus small, dense, positively charged center of an atom, contains protons and neutrons

Neutron neutral particle in the nucleus that has a mass nearly equal to the proton

Proton particle in the nucleus with a charge of 1+

Atomic Mass uses porous barriers to spate solids from liquids

Atomic Mass Unit uses different boiling points of substances to separate them

Atomic Number number of protons in an atom of an element

Isotope two atoms are isotopes if they have the same number of protons but different numbers of neutrons

Mass Number sum of the number of protons and number of neutrons.\

Chapter 4 section 4 doesn’t seem like one that we studied. However, it may be a good idea to look it it.

Notes

Democritus- 400BC

Though that the world was made of tiny individual particles

Atoms are solid, indivisible, and indestructible

Different kinds of atoms have different sizes and shapes

It is these size and shapes that give atoms different properties

Aristotle- 350BC

Rejected Democritus. He denied the existence of atoms and was so influential that this denial went largely

unchallenged for the next 2000 years or so.

Matter is made of earth fire, air, water.

Dalton- 1800

Matter is composed of extremely small particles called atoms

Atoms are indivisible and indestructible

All atoms of a given element are identical in mass and properties

Compounds are formed by a combination of two or more different kinds of atoms

A chemical reaction is a rearrangement of atoms

J.J. Thompson

Discovered electron

Robert Millikan

Calculated charge of electron by making an atom of oil drop and measure the charge it took to control its

descent.


29

The Structure of the Atom: Ch4 –cont.

The Plum Pudding model was commonly accepted, until Rutherford came along and did his gold foil experiment. He shot

electrons at thin gold foil, expecting them to go through. However, some actually bounced back, which led him to

discover the nucleus, which he said contained positively charged protons.

Soon after, James Chadwick discovered the neutron.

The atomic number is the number of protons, which will usually also be the

number of electrons. Because the amount of neutrons in atoms of the

same element may vary, the atomic mass is actually a weighted (no pun

intended) average of the different weights.

We can also write this in formation in the format:

Writing an element as carbon-14 indicates that the atomic mass is

14. Knowing this, you can subtract to find that there are 8 neutrons

In this isotope.

The atomic mass unit is about the weight of a

Proton or neutron, but it was calculated using

A carbon-12 atoms.

In this format, 63 is the amount of

protons plus neutrons. 29 is the amount

of protons. This tells you that there are

29 protons, 29 electrons, and 34

neutrons.

If you’re confused, I did 63-29= 34


h mν

Electrons in Atoms: Ch 5

Key Terms:

Amplitude wave’s height from origin to crest/trough

Wavelength shortest distance between equivalent points on continuous wave; crest-crest/trough-trough

Frequency number of waves that pass a given point per second

Atomic emission spectrum set of frequencies of the electromagnetic waves emitted by atoms of that element

Electromagnetic radiation form of energy that exhibits wavelike behavior as it travels through space

Electromagnetic spectrum includes all forms of electromagnetic radiation

Photoelectric effect the strange effect that not all light can eject electrons from a metal

Photon massless particle that carries a quantum of energy

Planck’s constant shows that the energy of radiating increases as the radiation frequency increases

Quantum minimum amount of energy that can be gained or lost

Atomic orbital 3D region around the nucleus which describes the electron’s probable location

Energy sublevel “slots” for pairs of electrons, as n increases, there are more sublevels Book pg 153

Ground state lowest allowable energy state of an atom is called its ground state

Principal energy level each major energy level of n Book pg 153

Principal quantum number indicates the relative size and energy of atomic orbitals Book pg 153

Quantum mechanical model model in which electrons are treated as waves

Quantum number the value of n that specifies orbital Book pg 153

Electron configuration arrangement of electrons in an atom

Electron-dot structure elements symbol surrounded by a number of dots equal to that elements valence electrons

Valence electron electrons in the outermost orbitals

Formulas:

c = 3.00 x 108 (m/s) speed of light

h = 6.626 x 1034 (j·s) Planck’s constant

λ (m)* wavelength

ν (Hz, 1/s)* frequency Frequency is measured in Hertz. 1 Hertz is 1 wave per second.

c = λν Electromagnetic Wave Relationship

Ephoton = hν Energy of a Photon

λ= Particle Electromagnetic-Wage Relationship* m is the mass of the particle.

Wavelength has to be in meters for formulas, so

don’t forget to convert!

Chemistry Honors Study Guide| /Notes 10

Electrons in Atoms: Ch 5

Notes

The black line is called the origin. In real life, this

wave is moving. The amount of times each crest

(called peak in this picture) passes you, that is

one Hz.

Make sure to understand that high energy waves have high

frequencies (they are fast) and short wavelengths. Low energy

waves have low wavelengths and low frequencies (they are

slow)

De Broglie equation

All moving particles have wave characteristics.

Heisenberg Uncertainty Principle

Its impossible to know both the location and speed of a particle. (the more you know about one, the less you know

about the other)

Hund’s Rule

Single electrons with the same spin must occupy each equal energy orbital before additional electrons with opposite

spins can occupy the same orbitals.

Aufbau Principle

Each electron occupies the lowest energy orbital available.

Pauli Exclusion Principle

A maximum of two electrons can occupy a single atomic orbital.

Hey guys, I didn’t cover a couple things here because the book does a

great job already. These topics are: Electron Configurations, Orbital

Diagrams, and Dot Diagrams.


The Periodic Table and Periodic Law: Ch 6

Key Terms:

Periodic law there is a periodic repetition of chemical and physical properties of elements

Group columns on the period table, also called families. These share traits.

Period rows on the periodic table

Representative Element elements in groups 1,2,13,18

Transition element elements in groups 3-12

Metal shiny, solid at room temperature, ductile, malleable

Alkali metal group 1, extremely reactive

Alkaline earth metal group 2, highly reactive

Transition metal an element in groups 3-12

Inner transition metal t he lanthanide and actinide series that is set off usually at the bottom of the graph

Lanthanide series f block elements from period 6 (starting with lanthanum)

Actinide Series f block elements from period 7 (starting with actinium)

Nonmetal elements that are generally gasses or brittle, dull-looking solids

Halogen group 17, very reactive

Noble gas group 18, extremely unreactive

Metalloid elements that have physical and chemical properties of both metals and nonmetals

Electronegativity ability of an element to attract electrons to it.

Ion atom or bonded group of atoms that has a positive or negative charge

Ionization energy energy required to remove electron from gaseous atom

Octet rule atoms tend to gain, lose, or share electrons in order to acquire a full set of 8 valence electrons

Notes Lavoiser, Newlands, Meyer and Mendeleev all worked to create the periodic table. Mendeleev is generally given

the most credit for organizing the elements into a series of rows of elements with similar characteristics. Mosley was able to then add more and slightly reorganize the table.


The Periodic Table and Periodic Law: Ch 6

Notes

Periodic Trends

As you go from top to bottom on the

periodic table, obviously your element will

be bigger. However, when you go from left

to right, since the electrons in that row are

in the same energy level regardless, the

more protons in the atom, the harder they

pull the electrons in close.

Also, when comparing ions, if the amount of

electrons are equal, then whichever has less

protons will be larger. If the amount of

protons are the same, then the one with

more electrons is bigger.

The easiest way for me to remember this

one is that Fluorine is the most

electronegative element. As you get further

away from it, you get less electronegative.

Chemistry Honors Study Guide| /Notes 13

Ionic Compounds and Metals: Ch 7

Key Terms:

Anion negatively charged ion

Cation positively charged ion (Why is Willy the wildcat a cation...Because he’s pawsative :)

Chemical Bond force that holds two atoms together

Crystal Lattice 3D arrangement of particles in which each positive ion is surrounded by negative ions

Electrolyte ionic compound whose aqueous solution conducts an electric current

Ionic bond bond in which one atom “takes” the others electron, often between a metal and nonmetal; NaCl

Ionic compound compound with ionic bond

Lattice energy energy required to separate 1 mol of the ions of an ionic compound

Formula unit chemical formula for an ionic compound (similar to a molecule)

Monatomic ion one atom ion; Mg2+

, Br -

Oxidation number charge of a monatomic ion

Oxyanion polyatomic ion composed of an element, usually a nonmetal, bonded to one/more oxygen atoms

Polyatiomic ion ions made up of more than one atom

Alloy mixture of elements that has metallic properties

Delocalized electron small, dense, positively charged center of an atom, contains protons and neutrons

Electron sea model all the metal atoms in a metallic solid lend their valence electrons to form a “sea” of electrons

Metallic bond attraction of a metallic a metallic cation for delocalized electrons

Notes

To the left is a picture of the electron sea

model.

Naming Ionic Compounds has been placed

at the “Chemistry Skills” section at the end

of this document.


Covalent Bonding: Ch 8

Key Terms:

Covalent bond bond in which atoms share valence electrons

Molecule formed when two or more atoms bond covalently

Lewis Structure Take a dot diagram, replace bonds with lines (page 242-244)

Sigma Bond single covalent bond (σ)

Pi Bond double covalent bond (π)

Endothermic reaction requires energy, generally becoming cold

Exothermic reaction releases energy, generally in the form of heat

Oxyacid an acid that contains both a hydrogen atom and an oxyanion

Structural formula similar to Lewis structure, but without dots. Does not show lone pairs.

Resonance condition that occurs when there is more than one way to draw the Lewis structure

Coordinate covalent bond forms when one atom donates both the electrons to be shared with an atom or ion that

needs two electrons to form a stable electron arrangement.

VSEPR model Valence Shell Electron Pair Repulsion, used to determine 3D shape of a molecule

Resonance a process that carbon atoms undergo in which atomic orbitals mix and form new shapes

Polar Covalent Bond polar covalent bond

Notes

Covalently bonded atoms.

Sigma bonds are single bonds, and are direct

between the two carbon atoms. However, when

atoms form double or even triple bonds, they have

to form them where there is not already one

present. So the pi bonds, which signify double/triple

bonds go around.



Notes

All the three figures at the top are valid. To show this, we

use dotted lines as it shows on the bottom. However, you

can ignore the numbers (1-,2/3-); they are unnecessary for

the purposes of Chem Honors.

Above is a little chart the differences between the Lewis Structure in 2D and the 3D representation. I believe we actually

need to know the differences between the different Tetrahedral figures, but don’t quote me on that.

Electronegativity Difference Bond Character

>1.7 Mostly ionic

0.4-1.7 Polar covalent

<0.4 Mostly covalent

0 Nonpolar covalent

This chart here is how you determine how to categorize the atoms. The teachers don’t expect you to memorize the

values, but think about it like this: the closer to each other two elements are, the more they will share electrons. This is

why metals and nonmetals form mostly ionic bonds with each other- they don’t share very well. Elements like oxygen

will form pretty much perfect nonpolar covalent bonds with themselves.



Key Terms:

Chemical reaction the process by which the atoms of one or more substances are rearranged into new ones

Reactants starting substances

Products substances formed

Chemical equation statement that uses chemical formation to show identities and amounts of substances

Coefficient number in front of product or reactant

Synthesis a chemical reaction in which two or more substance react to produce a single product

Combustion releases energy, generally in the form of heat

Decomposition an acid that contains both a hydrogen atom and an oxyanion

Single-replacement similar to Lewis structure, but without dots. Does not show lone pairs.

Double-replacement condition that occurs when

Precipitate solid produced during chemical reaction

Aqueous solution one or more solutes dissolved in solution; saltwater

Solute substance dissolved into the solvent; salt

Solvent substance solute is dissolved into; generally water

Complete ionic equation equation that shows all the particles in a solution as ions

Spectator ion ions that do not participate in a reaction

Net ionic equation ionic equations that include only particles that participate in reaction

Notes

I’m assuming you guys know how to balance a chemical equation…but just in case, check page 287 in your

book for a step by step guide.

Types of Reactions

Synthesis

Combustion

Decomposition

Single-replacement

Double-replacement

Chemistry Honors Study Guide| Total/Net Ionic equations 17

Total/Net Ionic equations

Example Expirement

Molecular Equation: CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) + CO2(g)

Total Ionic Equation: CaCO3(s) + 2H+(aq) + 2Cl

-(aq) Ca

2+(aq) + 2Cl

-(aq) + H2O(l) + CO2(g)

Net Ionic Equation: CaCO3(s) + 2H+(aq) Ca

2+(aq) + H2O(l) + CO2(g)

If a reactant or product is aqueous, split it up in your total ionic equation. If not aqueous (if its solid, gas, liquid) then

leave them as is. If there is a precipitate (a solid forms on the right side that wasn’t there on the left) then generally your

net Ionic equation will show the formation of that solid. Ignore the previous sentence if it confuses you.

Solubillity Rules

1. All common compounds of Group I and ammonium ions are soluble.

2. All nitrates, acetates, and chlorates are soluble.

3. All binary compounds of the halogens (other than F) with metals are soluble, except those of Ag,

Hg(I), and Pb. Pb halides are soluble in hot water.)

4. All sulfates are soluble, except those of barium, strontium, calcium, lead, silver, and mercury (I).

The latter three are slightly soluble.

5. Except for rule 1, carbonates, hydroxides, oxides, silicates, and phosphates are insoluble.

6. Sulfides are insoluble except for calcium, barium, strontium, magnesium, sodium, potassium,

and ammonium.

Here are some basic rules about solubility. Remember to check with the ones that your teacher gave you!


The Mole: Ch 10

Key Terms:

Mole SI base unit to measure amount of substance. A mol of something is 6.02x1023

of that thing.

Avogadro’s Number 6.02x1023

(originally calculated as the amount of carbon atoms in exactly 12 g of carbon-12)

Molar mass the mass in grams of one mole of any pure substance; the molar mass of water is 18g

Percent composition percent by mass of each element in a compound is the percent composition

Empirical Formula formula with the smallest whole-number ratio of the elements

Molecular Formula specifies the number of atoms of each element in one molecule of that substance

Hydrate compound that has a specific number of water molecules bound to its atoms

Notes If you have trouble understanding mols, think about it like this: if I had a mol of cars, I would have 6.02x1023

cars. If there was a mol of pages in this study guide, there would be 6.02x1023 pages.

If I had a mol of carbon, I would have 6.02x1023 atoms. If I had a mol of H2O, I would have 6.02x1023

molecules of water, 6.02x1023 atoms of oxygen. As for hydrogen, I actually have 2x6.02x1023, or 1.2x1024

atoms of hydrogen. Here’s why: for every 1 molecule of oxygen, you have 2 atoms of

hydrogen and 1 atom of oxygen. Hence the doubling.

To convert from grams to mols, divide the grams by the molar mass

To convert from mols to atoms/molecules, multiply the mols by 6.02x1023

To convert from mols to grams, multiply the grams by the molar mass

To convert from atoms/molecules to mols, divide the atoms by 6.02x1023

For percent composition and hydrates, my explanation

would not be any shorter than the book’s. Read pgs.342-

353. Sounds like a lot, but most of it is math. However, it is

important math :)


Stoichiometry: Ch 11

Key Terms:

Stoichiometry the study of quantitative relationships between the amounts of reactants and products

Mole Ratio ratio between the numbers of moles of any two of the substances in a balanced chemical equation

Limiting reactant limits amount of product created

Excess reactant reactants that are not completely used up in the reaction

Theoretical yield maximum amount of product that can be produced from a given amount of reactant

Actual yield amount of product poruduced when chemical reaction is carried out

Percent yield ration of the actual yield to the theoretical yield

Notes

Stochiometry

The thing about stoichiometry is that it is just a series of grams to mole conversions. However, its easy

to get confused by the steps.

Lets take the above equation. See how you need 2 groups of 2 hydrogen atoms and 1 group of 2 oxygen

atoms? That leads us to a total of 4 atoms of Hydrogen and 2 Oxygen. This is a 2:1 ratio. So if I had 1 mol

of O2, I would need 2 mol of H2. The coefficients in front of the letters is just a ratio of mols. So if I

had 2.5 mol O2, how many mol H2O would I make? 2x2.5= 5! I would make 5 mol H2O.

o A little trick for remembering how to use stoichiometric ratios goes as follows:

1) Make sure you have your reactants/product in moles

2) You would like to convert mol of one substance, lets call it A, to another, called B.

2) Take your moles of substance A, and divide it by the coefficient of substance A in the

equation.

3) Then, multiply this new number by the coefficient of the substance B.

4)The number you get is your mols of substance B

Example problems for this can be found on pages 375-377


Limiting Reactants

This is easy to understand if you think about it in other terms first.

Each car takes exactly 1 car body and 4 tires to make. In this example, we have 48 tires. 48/4=12. We have enough tires

for 12 cars. We also have 8 car bodies so we have enough to make 8 cars. Therefore, we can only make 8 cars, no matter

how many tires we have.

Lets take this problem: A 2.00 g sample of ammonia is mixed with 4.00 g of oxygen. Which is the limiting

reactant and how much excess reactant remains after the reaction has stopped?

First, write the chemical equation: 4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g)

Then, plug in the two values to see which substance makes the most of the product. Because the question did not

specify, we can just pick a product; lets use NO.

Above is the conversion the way that the book would do it. They calculated the mols of NO created twice, using the 2 g

of ammonia the first time and the 4 grams of oxygen the second, then converted it in grams. We have enough ammonia

for 3.53 g NO. We have enough oxygen for 3 g NO. Therefore, at most we can make 3 g NO. Oxygen is the limiting

reactant.

For the second part of the problem, we have to do some sneaky math. We know that oxygen is the limiting reactant; so

why not just convert mol of oxygen into grams of ammonia? Doing this conversion (divide by molar mass O2, divide by 5,

multiply by 4, multiply by molar mass NH3) gives you the amount of NH3 that was used. So if 2 g were given to us, and we

used 1.7, how much is left? .3 g NH3

Theoretical/Actual Yield

Its important to remember that when you do experiments in real life, results aren’t always perfect. The theoretical yield

is the amount that the math says you’ll get; the actual yield is what happens in a lab setting. Your percent yield is

therefore a percentage of how well your lab compared to a perfect one. For example, I plugged in my 4 grams of Oxygen

into my chemical formula (this is an imaginary formula, don’t look for it) and found my theoretical yield to be 10 grams

of Magnesium. However, when I did the lab, I only got 6 grams of magnesium. 6/10=.6….my percent yield was 60%.


States of Matter: Ch 12

Key Terms:

Elastic collision on in which no energy is lost; think if it like pool balls as opposed to playdough

Temperature measure of the average kinetic energy

Diffusion describes movement of one material through another

Effusion describes how a gas escapes through a small hole

Pressure force per unit area; kPa, atm, mmHg, torr

Barometer instrument that measures atmospheric pressure

Pascal force of one newton per square meter, unit for pressure

Atmosphere average air pressure at 0 Celsius and sea level, unit for pressure

Dispersion force weak forces that result from temporary shifts in electron density

Dipole-dipole force attractions between oppositely charged regions of polar molecules

Hydrogen bond dipole-dipole bond between hydrogen and N, O, or F. stronger than regular dipole-dipole

Viscosity measure of a resistance to flow (the difference between pouring water and syrup)

Surface tension energy required to increase the surface area by a given amount

Surfactant compounds that lower the surface tension of water

Crystalline solid solid whose atoms, ions, or molecules are arranged in an orderly geometric structure

Unit cell smallest arrangement of atoms in a crystal lattice that has the same structure as the whole

Allotrope an element that exists in different forms at the same state

Amorphous solid one in which the particles are not arranged in a regular, repeating pattern

Melting point temperature at which the forces holding a crystal lattice together break

Vaporization process of a liquid changing to a gas

Evaporation vaporization that occurs only at the surface of a liquid

Vapor pressure pressure exerted above the liquid

Boiling point temperature at which the vapor pressure of a liquid equals the atmospheric pressure

Freezing point temperature at which a liquid is converted into a crystalline solid

Condensation process by which a gas becomes a liquid

Deposition process by which a substance changes from a gas to a sold without being a liquid

Phase diagram graph of pressure

versus temperature, shown to the right

Triple point point on a phase

diagram at which all three

phases can coexist

Critical point point above which

a substance cannot exist as a liquid


Stoichiometry: Ch 12- cont.

Notes Kinetic-molecular theory

1. Gases consist of large numbers of molecules (or atoms, in the case of the noble gases) that are in

continuous, random motion

2. The volume of all the molecules of the gas is negligible compared to the total volume in which the gas is

contained

3. Attractive and repulsive forces between gas molecules is negligible

4. The average kinetic energy of the molecules does not change with time (as long as the temperature of

the gas remains constant). Energy can be transferred between molecules during collisions (but the

collisions are perfectly elastic)

5. The average kinetic energy of the molecules is proportional to absolute temperature. At any given

temperature, the molecules of all gases have the same average kinetic energy. In other words, if I have

two gas samples, both at the same temperature, then the average kinetic energy for the collection of gas

molecules in one sample is equal to the average kinetic energy for the collection of gas molecules in the

other sample.

Graham’s law of effusion

The two M’s are the molar masses of the two gasses.

Dalton’s law of partial pressures

PressureTotal = Pressure1 + Pressure2 ... Pressuren Dalton’s law of partial pressures basically states that if you know the total pressure, then you can subtract the different parts. For example, the total pressure on my head from a combination of nitrogen and oxygen is 1.5

atm. The nitrogen has a pressure of 1.25 atm. What is the pressure of Oxygen? 1.5-1.25= .25

Unit Number Equivalent to 1 atm

kPa 101.3

Atm 1

mmHg 760

Torr 760

Psi 14.7

Bar 1.01

(mm Hg is the same thing as Torr. We don’t use Psi or Bar very often.)

Dispersion forces are weaker than dipole-dipole forces which are weaker than hydrogen bonds. If you struggle

with understanding what dispersion forces are, remember that the electrons in an atom are moving around randomly.

Dispersion forces are the attractions that occur when it just so happens that the electrons in molecules are more

numerous on one side than another.

Only crystalline solids have a melting point. Amorphous solids actually have more of a range of temperatures

where they melt. Also,


temperature and pressure are………… directly proportional

pressure and volume are…………………..inversely proportional

volume and temperature are…………….directly proportional

Gases: Ch 13

Key Terms: Absolute zero zero on the Kelvin scale, at which temperature all molecular motion stops

Molar volume volume that 1 mole of a gas occupies at STP, 22.4 L

Avogadro’s principle equal volumes of gas at the same temperature and pressure contain equal numbers of

particles

Formulas:

Boyles Law

Charles Law

Gay-Lussac’s law

Combined Gas Law

Ideal gas law

Notes:

o As the temperature goes up, the pressure goes up

o As the temperature goes down, the pressure goes down

o As the temperature goes up, the volume goes up

o As the temperature goes down, the volume goes down

o As the volume goes down, the pressure goes up

o As the volume goes up, the pressure goes down

The ideal gas constant is represented by

the symbol R. Depending on the unit of

pressure, there are different values of R.

atm .0821

kPa 8.314

mmHg 62.4

Personally, I just remember the R value

for atm and just convert. But ,since we

get to use a cheat sheet, this would

probably be a good idea to write down.

Also, don’t forget to use Kelvin instead

of Celsius of Fahrenheit when

performing these calculations.

STP is 0 Celsius, 1 atm


Gases: Ch 13

Notes: Most gasses are more or less ideal at most conditions. However, they are furthest from being ideal gasses at

high pressures and low temperatures.

Polar gasses (water vapor) do not behave ideally

Gas Stoichiometry

Gas stoichiometry is not that much different than regular stoichiometry. The main difference here is that you

often have to use conversions. The key thing that keeps you from gas stoichiometry being just regular stoichiometry is

the fact that you often are given gas as a volume. Use the gas laws to find out moles, then its back to basics.

This is a shortcut, but you have to make sure that the problems you are given tell you to assume that pressure

and temperature remain constant. This means that you can actually just take the volume in L that you were given and

treat them like mols, without converting. Example on page 461.

I’ll admit, Gas stoichiometry isn’t that easy. It is highly advisable to check out the

information on pages 461 to 464 in your book.



Significant Digits

Digits from 1-9 are always significant.

Zeros between two other significant digits are always significant One or more additional zeros to the right of both the decimal place and another significant digit are significant.

Zeros used solely for spacing the decimal point (placeholders) are not significant.

Examples of Significant Digits

EXAMPLES # OF SIG. DIG. COMMENT

453 kg 3 All non-zero digits are always

significant.

5057 L 4 Zeros between 2 significant

digits are significant.

5.00 3

Additional zeros to the right of

decimal and a significant

digits are significant.

0.007 1 Placeholders are not

significant.

Adding and Subtracting

RULE: When adding or subtracting your answer can only show as many decimal places as the

measurement having the fewest number of decimal places.

Exampled: When we add 3.76 g + 14.83 g + 2.1 g = 20.69 g

We look to the original problem to see the number of decimal places shown in each of the original

measurements. 2.1 shows the least number of decimal places. We must round our answer, 20.69, to one decimal

place (the tenth place). Our final answer is 20.7 g


Significant Digits –cont.

Multiplying and Dividing

RULE: When multiplying or dividing, your answer may only show as many significant digits as the

multiplied or divided measurement showing the least number of significant digits.

Example: When multiplying 22.37 cm x 3.10 cm x 85.75 cm = 5946.50525 cm3

We look to the original problem and check the number of significant digits in each of the original

measurements:

22.37 shows 4 significant digits.



Our answer can only show 3 significant digits because that is the least number of significant digits in the

original problem.

5946.50525 shows 9 significant digits, we must round to the tens place in order to show only 3 significant

digits. Our final answer becomes 5950 cm3.


Naming Chemical Compounds

Binary Molecular Compounds

Binary molecular compounds are composed of two types of nonmetals. The nonmetals are normally ordered

with the element leftmost on the periodic table first. If both elements are in the same column, then the element

lower on the periodic table is first. The order is the same for the formula and the name.

Because the nonmetals can combine in many ways, prefixes are used to express how many atoms of each

element are in the atom. Memorize the following list of prefixes.

1 mono* 2 di 3 tri 4 tetra 5 penta

6 hexa 7 hepta 8 octa 9 nona 10 deca

The prefix for one is starred because its use is optional for the second element. If there is only one atom of the

first element, no prefix is used.

Examples:

Carbon Monoxide CO

Oxygen Difluoride OF2

Tetranitrogen Decaoxide N4O10

Arsenic Tribromide AsBr3

Boron Trichloride BCl3

Iodine Heptafluoride IF7


Binary ionic compounds may contain more that two elements but are binary because they contain two ions.

First, take your cation (it will usually be a metal from the left). Make no changes; this will be the first part.

Then, take your negative part (usually an element from the right side) add it on to the end. Drop the ending, add

an –ide. If your working with a polyatomic ion, just stick it in as is. They are already “conjugated.”

Remember, Alkali metals always have a charge of +1. Alkali earth metals always have a +2 charge. In addition

to these two groups, aluminum is always Al3+, zinc is Zn2+, and silver is Ag+.

Examples:

sodium carbonate Na2CO3

cobalt(II) nitrate Co(NO3)2

tin(IV) sulfide SnS2



Since acids are substances that release H+ in water, it is traditional to write the hydrogen atom first in the

formula. The names of these acids are based on the anion the acid came from. (Hydrogen acts as a cation, H+.

Although acids are molecular compounds, they react with water to form ions.) If the anion has an ate ending,

the ate is changed to ic and the word acid added. If the anion has an ite ending, the ite is changed to ous and is

followed by acid.

Examples:

hydroiodic acid HI (that’s H with a capital i after it)

carbonic acid H2CO3

sulfurous acid H2SO3

perchloric acid HClO4

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