Star Test Review

CHEMISTRY

BREAKDOWN OF QUESTIONS

6 Questions on the Periodic Table•To relate the position of an element in the periodic table to its atomic number and atomic mass

•Use the periodic table to identify metals, semimetals, non-metals, and halogens

•Identify alkali metals, alkaline earth metals and transition metals, trends in ionization energy, electronegativity, and the relative sizes of ions and atoms…(2 Questions on this…)

•Use the periodic table to determine the number of electrons available for bonding

•The nucleus of the atom is much smaller than the atom yet contains most of its mass.

The Modern Model

electron

neutron

proton

(not to scale) Chadwick’s neutrons

Rutherford’s space and nucleus

Dalton’s atom

Bohr’s energy levels

Thompson’s electrons

Elements• 110 known elements– 92 of which are naturally

occurring.

• Each has an atomic symbol.

• Atomic number– is number of protons

• Atomic mass– is the total mass of the protons

plus the neutrons.

OOXYGEN

8

15.9994

Notice that the atomic mass is not a round number, even though protons and neutrons each have a mass of 1. This is due to natural abundance.

Energy Levels (n)• The electrons exist in energy

levels or shells.

• The first energy shell can hold only 2 electrons.– Hydrogen and Helium in their

ground state have electrons that occupy this shell.

• The second shell can hold 8 electrons.

• The third can hold 18 electrons.

2 8

32

18

Shells

All shells after three can hold 32 electrons.

Ion e- configurations

• Ions (elements with more/less electrons) also have electron configurations.

• Consider Sulfur (S):• What if sulfur gained an electron?

• Consider Calcium (Ca):

• What if calcium lost two electrons?

4233]Ne[S ps

5233]Ne[-S ps

262622 433221Ca spspss

626222 33221Ca pspss

• Question:– Why do the atomic radii of atoms decrease as

electrons are added to the atom, as you move from left to right across a period?

• electrostatic attraction– attraction between the electrons (-) in the shells and

the protons(+) in the nucleus – pulls the electrons in

This is what we call a periodic trend

The Periodic Table• The Periodic Table– a collection of all the known elements into a model

that groups elements with similar properties.

• Groups– Vertical columns of elements with similar properties.

• Periods– Horizontal rows of elements with atomic mass and

similar electron configurations.

Periodic Table History• Dmitri Mendeleev– Russian chemist created who ordered the known

elements according to properties. (Gaps?)

• Henry Moseley– arranged the elements according to atomic number

(# of protons).

– This is the system we use today.

• Periodic Law– chemical and physical properties of elements are

periodic functions of their atomic numbers.– The elements in the periodic table are arranged

according to Periodic Law– Periodic Law shows certain trends in the properties

of elements

c 1869

c 1911

Periodic Trends – Atomic Radii• As electrons are added to the outside of atoms,

in the same period, the atom’s radius decreases.

• As new shells are added, radius increases.

Smaller from left to right

Periodic Trends – Ionization Energy• Ionization Energy - the energy required to strip an

electron from an atom.• As more electrons are added to a shell, it’s more

difficult to remove them. (More protons pulling inward)• Easier to remove electrons from larger atoms.

Larger from left to right

eAenergyA

Period Trends – Electronegativity

• Electronegativity (electron affinity)– an atom’s ability to attract electrons – Negative electron affinity = atom wants e-.– Decreases down a group

Larger from left to right

Ionic Radii• Recall: + - attraction

determines the atom’s radius.

• An electron is added to a nonmetal atom :– Anion is formed.– Anions are larger than their neutral

atom

• An electrons is removed from a metal atom:– Cation is formed.– Cations are smaller than their

neutral atoms

)( anioneatom

Cl Cl-

Na Na+

)( cationeatom

Groups and their Properties• Recall:

– elements in the same group have similar properties due to similar electron configurations.

• Learn the following group-families and their basic chemical and physical properties:– Alkali Metals– Alkaline-Earth Metals– Transition Metals– Main-Block Elements– Noble Gasses– Rare-Earth Elements

Group 1 (+1)Alkali Metals (s)

• soft, highly reactive metals.• Lustrous

– will reflect light, but these elements quickly lose their sheen when exposed to the air.

• Electrically Conductive– able to pass a charge through the material. These

elements are often found in lights, batteries, and electrolytes.

• have low melting points• low density.

1#sconfigeending

Group 2 (+2)Alkaline-Earth Metals (s)

• Properties are similar to group 1 elements, but are:

• Harder• Less reactive than Group 1 elements.

– (These elements are still very reactive.)

• Lustrous• Electrically Conductive• Higher melting points than Group 1 metals.• More dense than Group 1 metals.

2#sconfigeending

Groups 3-12 (various)Transition Metals (d)

• This is where we find most metals, including the coinage metals.

• Lustrous• Electrically Conductive• Malleable

– able to be shaped and formed, and hold that shape.• Ductile

– able to be drawn into wires• Very hard• Very dense

• High melting points ##dconfigeending

Group 13-17 (+3-1)Main-Block Elements (p)

• The most varied elements.

– Liquids, gasses, and solids can be found in this group. With widely varied properties

• Includes Metalloids

– elements having properties of both metals and non metals.

• Most elements necessary to living things are found in this section.

• Includes Halogens

– Group 17 gasses and liquids F, Cl, Br, I, At

– are very reactive due to very high electron affinities.

## pconfigeending

Group 18 (0)Noble Gases (p)

• NearlyNearly unreactive.• All have filled octets.• Near zero electron affinity• Very high ionization energies.• Noble gasses make up a trace amount of our

atmosphere– are mined from pockets of gases in the oceans.

• When electrically charged:– noble gases produce brilliant plasmas, often used in

signs.

6# pconfigeending

f – Group (various)Rare earth metals (f)

• Very heavy, dense (large nuclei)• Most are radioactiveradioactive.• Lanthanides

– The first row, starting with lanthanum (57La)

– (4f elements)• Actinides

– The second row, starting with actinium (89Ac)

– (5f elements)• Transuranium elements

– All elements after Uranium 92U (93Np on) are artificial.

#54 forconfigeending

Chemical Bonds---- 7 Questions •Atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds. (2 Questions)

•Know chemical bonds between atoms in molecules such as H2, CH4, NH3, H2CCH2, N2, Cl2 and many large biological molecules are covalent.

•Salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by electrostatic attraction

•Atoms and molecules in liquids move in a random pattern relative to one another because the intermolecular forces are too weak to hold the atoms or molecules in a solid form

•Know how to draw Lewis dot structures. (2 Questions)

Chemical Bonding and Molecular Compounds

Ionic bondIonic bond:

Electrons are given or taken, occurs between metal and nonmetal (NaCl)

Covalent bondCovalent bond:

Electrons are shared, occurs between two or more nonmetals (CO2)

Ions• Ion– atom that has gained or lost one or

more electrons.

• Octet Rule– all atoms want zero or eight

electrons in their outer shell.

• Metal and nonmetal bond:– each atom seeks to gain electrons

or lose them. For instance:

• Cl wants to gain one electron• Na wants to lose one electron• Notice the sizes of the ions

Cl Cl-

Na Na+

Positive ions are called “cations”

Negative ions are called “anions”

+e-

-e-

Ionic attraction• Oppositely charged ions

attract– Like opposite poles of a

magnet attract

• Positively charged Sodium Ion and negatively charged Chlorine Ion are attracted.

• They remain ions, but stick together in a lattice– (3D grid pattern) as other ions

join them.

Cl- Na+

Cl- Na+

Cl-Na+

Na+

Cl-

Formula Units• Na+ and Cl- combine

to form NaCl.

• NaCl is one formula unit:– the smallest unit that has the

correct formula for our compound.

• NaCl is also our compound’s empirical formula:– the smallest ratio of atoms that

make up our compound.

• Although there are several Na+ ions and Cl- ions in the lattice, the formula unit is still just NaCl.

Cl-Na+

Metallic Bonding• In metallic bonding

– electrons are shared and flow between metal atoms in an “electron sea.”

– This is an extremely strong bond• This is why metals:• Have luster

– (reflect light)• Are malleable

– (able to formed and hold that form)• Are ductile

– (able to formed into wires)• Conduct electricity

– (pass an electrical charge)• Conduct heat

– (pass heat easily)

d-orbital electrons

Covalent Bonding

• Covalent bonds are bonds formed when– atoms share electrons to complete their octet.– This is a very strong bond.

• Atoms on the upper-right of the periodic table (nonmetals) covalently bond with each other. (except nobles)

• Molecular compounds– covalently bonded compounds.

• Diatomic molecules are covalently bonded.

– I2, O2, F2, Cl2, H2, Br2, N2.

Covalent Bond Energy• Bond length and bond energy are inversely

related.– That is, a smaller bond has a greater energy, and

will release more energy when broken.• Atoms will bond at a distance that is most stable

– (lowest potential energy)

High potential energy = unstable

Low potential energy = more stable

Electronegativity• Electronegativity:

– Atom’s pull for electrons in a covalent bond

• Electrons are sometimes not shared equally among atoms.

• Polar covalent– Electrons are shared, but one

atom has a higher Electronegativity

– Like in CO.

• Nonpolar covalent– Electrons are shared evenly– Like in O2. O O

OC

Cl-Na+

Molecular notation - Lewis Dot• Lewis Dot structures

– tool used to draw molecules in 2D with stick diagrams.

• Steps:– Determine each atom’s number of valence

electrons– Number of valence electrons = number of dots– Molecules share dots until each atoms is

surrounded by 8 dots, representing a completed octet.

– Hydrogen gets only 2 dots on one side– Pairs of dots between atoms are changed to sticks

– representing covalent bonds.– Any pairs of dots that are not shared between

atoms are called unshared pairs.

Lewis Dot Molecular Model

• Consider Methane, CH4

• The atoms are C and H

• Add each atom’s valence electrons as dots, clockwise.

Lewis Dot Molecular Model

• Consider Methane, CH4

• The atoms are C and H

• Add each atom’s valence electrons as dots, clockwise.

• Combine the dots to create octets, carbon belongs in the center.

Lewis Dot Molecular Model

• Consider Methane, CH4

• The atoms are C and H• Add each atom’s valence

electrons as dots, clockwise.

• Combine the dots to create octets, carbon belongs in the center.

• Replace pairs of dots with sticks, representing bonds.

Lewis-Dot Practice:• Try the following on your own:

– Determine each atom’s number of valence electrons – Arrange until all atoms except H have 8 valence electrons.

(H has 2)– Replace ( : ) with sticks– (hint-carbon is always in the middle)

OH 2 3NHHCl 62HC OHHC 52

Conservation of Matter and Stoichiometry ---- 10 Questions

• Know how to describe chemical reactions by writing balanced equations. ( 2 Questions)

•Know the quantity one mole is set by defining one mole of carbon 12 atoms to have a mass of exactly 12 grams.

•know one mole equals 6.02 x 1023 particles (atoms or molecules).

•know how to determine the molar mass of a molecule from its chemical formula and a table of atomic masses and how to convert the mass of a molecular substance to moles, number of particles, or volume of gas at standard temperature and pressure. ( 3 Questions)

•Know how to calculate the masses of reactants and products in a chemical reaction from the mass of one of the reactants or products and the relevant atomic masses. (3 Questions)

Chemical ReactionsProducts and Reactants

• We write the compounds to react on the left and the compounds produced on the right.

• Remember, reactants react to produce products

producedcompoundsreacttocompounds

productsreactants

Read as “yields”

Symbols and Notation• (s) – compound is a solid• (l) – cmpd is a liquid• (g) – gas• (aq) – aqueous (dissolved in water, exists

as ions)• ↓ - a precipitate has formed

– (solid falling out of the reaction)

• ↑ - a gas is evolved– (bubbling out of the reaction)

• Various things can be put over the “yields” sign (→) to indicate special reaction conditions.

)(2 sPbI

)(2 lOH

)(4 gCH

)(aqKI

2PbI

2HCcatalyst00

Chemical Equations• We express a chemical reaction with a chemical

equation.• This shows relative number of products and reactants

required to satisfy the Law of Conservation of Mass.• Our lab reaction:

• Must be written:

• This balances the equation – both the reactants and products are equal.

)(3)(2)()(23)( aqsaqaq KNOPbIKINOPb

)(32)()(23 22)( aqaqaq KNOPbIKINOPb

Diatomic Molecules

• Some elements don’t exist alone, but must form a pair with itself.

• These are diatomic molecules.

• In chemical reaction, these elements must be written as a X2.

OHH 202

DiatomicMolecules

Hydrogen = H2

Nitrogen = N2

Oxygen = O2

Fluorine = F2

Chlorine = Cl2Bromine = Br2

Iodine = I2

222 202 OHH

Rules for Balancing Chemical Rxns

• Never Change Subscripts• Balance Groups First• Balance H2O, O2, and H2 last.

• We balance by adding coefficients.

• Two reactant oxygens• Four product hydrogens• Properly balanced

42SOH

OHOH 222

OHOH 222 22

Mole Ratios• Mole ratios:– how many moles of products are produced

with given a number of moles of reactants.

• Here, the mole ratio is 1:1:2 (1H2 : 1Cl2 : 2 HCl)• This means two moles of HCl will be produced

when one mole of H2 and one mole of Cl2 react.

Balancing Practice• Balance the following reactions on your own:

OHCOOOHC 2226126

Respiration

OHCOOHC 22222 Acetylene torch

FeOAlOFeAl 3232Thermite

322 OFeOFe Rust

66

6

2 5 4 2

2 2

24 3

+ Energy

+ Energy

+ Energy

Energy In and Out

• Recall that a reaction can be Exothermic (releasing energy) or

• Endothermic (absorbing energy)• Some exothermic reactions need a little

energy to get going, but once going, will give off more energy. (energy barrier)

• Activation Energy – energy needed to get the reaction

going. Exothermic

The Mole• The “mole” represents a number of

things….like a dozen.• How many things is a mole?• 6.022137 x 1023…but we will use 6.02

x1023.• This is Avogadro’s number

– named for a lawyer, Amadoe Avogadro, that studied molecular gasses (diatomic) as a hobby.

• When you have three moles of atoms, you have (3 x 6.02x1023 =) 1.81x1024 atoms total.

molesmolggrams 0.5/23115

Molar Mass• Molar mass– expressed in grams per mole (g/mol) is the mass of

one mole of a substance.

• The mole is the link between the very small (atoms) and the macro (grams)

• The average atomic mass of carbon is 12.01. What is the mass of a mole of carbon atoms?

• Sodium has an atomic mass of 23 g/mol. How many moles do you have in 115 grams?

12.01 grams!

5.0 moles!

Mole-Mass Conversions• Complete the following mole-to-mass conversions:

• Mass in grams of 2.25 moles of iron, Fe?

• 126 grams Fe

• Mass in grams of 0.375 moles of potassium, K?

• 14.7 grams K

• Number of moles in 5.00 grams of calcium, Ca?

• 0.125 moles Ca

• Number of moles in 3.60x10-10 grams of gold, Au?

• 1.83x10-12 mol Au

Use your periodic table to find molar mass

Mole-Atoms conversions• Mole = 6.02x1023 things, how many atoms are in:

• 3.0 moles of silver, Ag?

• 0.010 moles of copper, Cu?

• How many moles do you have in:

• 2.4x1024 atoms of helium, He?

• 3.0x1023 atoms of lithium, Li?

• How many moles do you have in 127.1 grams of copper?

• How many atoms in 127.1 grams of copper?

atomsxxx 2423 108.11002.60.3

atomsx 211002.6

molesx

x4

1002.6

104.223

24

moles5.0

molesmolggrams 2/55.63/1.127

atomsxxmolesmolggrams 2423 102.11002.62/55.63/1.127

Phew!

Gases and Their Properties --- 6 Questions

•Know the random motion of molecules and their collisions with a surface create the observable pressure on that surface.

•Students know the random motion of molecules explains the diffusion of gases.

•Know how to apply the gas laws to relations between the pressure, temperature, and volume of any amount of an ideal gas or any mixture of ideal gases. ( 2 Questions)

•Know the values and meanings of standard temperature and pressure (STP).

•Know there is no temperature lower than 0 Kelvin.

•Know how to convert between the Celsius and Kelvin temperature scales.

Properties of Gases

• Gases have very low density– particles are spaced far apart.

• Gases are compressible.– Extreme pressures-gases will compress until they

become liquids (or solids, CO2).

• At 1 atm and 0oC temperature (STP),– most gases have a standard molar volume of about

22.4 liters.– One mole of gas at STP has a volume of 22.4 liters.

About 1000 times less dense than solids!

Particle Theory of Matter• Recall: The particle theory of matter states

– all matter is made up of particles (atoms) in random and constant motion.

– Particles are continually bouncing off each other (colliding).

• Adding heat to a system– increases the temperature …– Temperature = measure of the average kinetic

energy of the particles.

• Increasing the pressure of a gas– increases the density of the gas - the number of

particles in a given space.

Boyle’s Law• A young, adventurous, British aristocrat named

Robert Boyle found that– when temperature is kept constant, volume varies

inversely proportional with pressure. That is:

• P V = k (constant)

• We tend to write Boyle’s Law as the volumes and pressures under two conditions:

• P V = P V

c 1660’s

Charles’ Law• French chemist,

Jacque Charles, showed that at constant pressure,– temperature and volume

varied proportionally. That is…

• V / T=k (k = some constant #)

• We tend to write Charles’ Law as the volumes and temperatures under two conditions:

c 1780’s

Acids and Bases --- 5 Questions

•Know the observable properties of acids, bases, and salt solutions. ( 2 Questions)

• Students know acids are hydrogen-ion-donating and bases are hydrogen-ion-accepting substances.

•Students know strong acids and bases fully dissociate and weak acids and bases partially dissociate.

•Students know how to use the pH scale to characterize acid and base solutions.

Properties of Acids

• Acids…• Taste sour (but don't taste them in lab!!) • Form Electrolytes (ions are in solution)• React with bases to form salts and water. • Turn Blue Litmus Paper to Red• Are generally thought of as being corrosive.

Strong acids will react with many metals, dissolving the metal atoms into ions.

• The hydronium ion (H3O+) is the “acid ion”

)()(32)( aqaqaq ClOHOHHCl

Properties of Bases

• Bases…• Have a slippery feel• Have a bitter taste (but don't taste them in lab!!) • Form Electrolytes (ions are in solution) • React with acids to form salts and water • Turn Red Litmus Paper to Blue• Although frequently less reactive than acids in the

presence of metals, strong bases pose a greater threat to biological material due to their hydrophilic nature.

• Hydroxide (OH-) “base ion”

)()()( aqaqaq OHNaNaOH

The pH scale• The “p” in pH means “negative log” (-log)

• The “H” in pH means hydronium ion [H3O+] concentration.

• so…pH means the negative log of the hydronium ion concentration -log [H3O+]

• Solutions with low pH have high H3O+ concentration. “Acids”

• Solutions with high pH have low H3O+ concentration. “Bases”

• When a solution increases in pH, does it become more acidic, or more basic?

The pH scale• Solutions with low pH have high H3O+ concentration. “Acids”

• Solutions with high pH have low H3O+ concentration. “Bases”

Solutions--- 3 Questions

• Students know the definitions of solute and solvent.

•Students know how to describe the dissolving process at the molecular level by using the concept of random molecular motion.

•Students know temperature, pressure, and surface area affect the dissolving process.

•Students know how to calculate the concentration of a solute in terms of grams per liter, molarity, parts per million, and percent composition.

Parts of a solution

• The dissolving medium is the solvent (what does the dissolving…the dissolver)

• The dissolved substance is the solute (what gets dissolved…the dissolvey)

• The solute and solvent together form the solution.

• Solvents and solutes can be any phase.

solution

Electrolytes

• Electrolytes– Solutions that conduct electricity.

• Ionic solutions are electrolytes.• Covalent solutions are nonelectrolytes.• Is saltwater (NaCl in water) an electrolyte?

• Is sugar water (C6H12O6 in water) an electrolyte?

• Conductivity tester– can tell us if a solution is an electrolyte, and

sometimes, how strong an electrolyte is.

Solubility• Solubility– The extent to which a solute will dissolve in a

solvent. (how much solute will dissolve)

• High solubility– large amounts of solute will dissolve in a solvent

• Low solubility– only small amounts of solute will dissolve

• Increasing temperature increases the solubility of solids in liquids

• Increasing temperature decreases the solubility of gasses in liquids! …

Solid-Liquid and Gas-Liquid solubility with temperature

Gasses in liquids• In addition to cold

temperatures, high pressures increase solubility of gasses in liquids.

• Henry’s Law:– solubility of a gas in

a liquid increases with increasing pressure of that gas above the liquid.

Like Dissolves Like!• Some solvents are polar, like magnets, having partial

negative and partial positive ends. (H2O)

• Other solvents are nonpolar, having no “+” “-” poles• Polar solutes tend to dissolve well in polar solvents…• Nonpolar solutes tend to dissolve well into nonpolar

solvents.• “Like dissolves like”• Water is very polar. Does it dissolve polar substances

or non polar substance?

Saturation

• Saturated Solution– solution has as much solute as it will allow (equal to

solubility)• Unsaturated Solution

– more solute can dissolve into solution (less than solubility)

• Supersaturated Solution– too much solute in solution-some will fall out (more

than solubility)• We express the quantitative amount of solute in a

solution with concentration …

Concentration - Molarity

• Concentration– the quantitative amount of solute present in a

solution

• Molarity (M) – moles/liter– number of moles solute in liters of solution

Try these Molarity questions

• What is the concentration [in Molarity] when 3 moles of NaCl are dissolved in 2 Liters of water?

• How much (in liters) of a 0.1 M solution do you need to get 2 moles of solute?

• How many moles of NaOH are present in 300mL of a 1M solution?

• How many grams of HCl are found in 100mL of a 2M solution?

1.5 M “molar”

20 L

.3 moles 7.2 grams

Solution Preparation• By solid dissolving:• 1. calculate how many grams are needed to create our

volume of our desired molarity solution• 2. weigh out that mass, and add it to a volumetric

flask• 3. add some water and allow to dissolve• 4. add water to the desired volume

• By dilution of a standard solution:• 1. use the relationship M1V1=M2V2

• 2. calculate volume of standard molarity solution to use to get desired volume of desired molarity solution.

Strong/Weak Electrolytes

• Recall that a solid compound made up of a cation and anion is called a salt.

• Salts that dissolve completely into their ions when put in water dissociate completely.

• Salts that dissociate completely form strong electrolytes – solutions that conduct electricity well.

• Some salts only partially dissociate, forming weak electrolytes – solutions that conduct electricity, but do so poorly.

H+ / OH- Ions – (Acids and Bases)

• When a H+ ion is released into solution, a H3O+ ion is produced, called Hydronium ion.

• When a OH- ion is produced, we call this a Hydroxide ion.

• Hydronium (H3O+) and Hydroxide (OH-) are the fundamental ions involved in acid/base chem.

• Acids that dissociate completely, releasing H+ ions form strong electrolytes.

• Bases that dissociate completely releasing OH- ions form strong electrolytes.

Chemical Thermodynamics --- 5 Questions

• Students know how to describe temperature and heat flow in terms of the motion of molecules (or atoms).

•Students know chemical processes can either release (exothermic) or absorb (endothermic) thermal energy.

•Students know energy is released when a material condenses or freezes and is absorbed when a material evaporates or melts.

•Students know how to solve problems involving heat flow and temperature changes, using known values of specific heat and latent heat of phase change. (2 Questions)

Energy

Thermochemistry:Causes of Change in Systems

Heat, Energy, and Temperature changes

Standard unit of heat is the Joule, J

Standard unit of temperature is Kelvin, K

Heat vs Temperature• Heat

– measure of energy change in a system.

• Temperature– measure of the kinetic energy (movement) of the

particles in a system.

• Gaining or losing heat energy in a substance can change its temperature.

• Exothermic– System loses energy to surroundings

• Endothermic– System gains energy from surroundings

Specific Heat Capacity

• Specific Heat Capacity– measure of how a substance reacts to heat energy

changes.

• The symbol we use is cp.– The “p” stands for constant pressure while heat is added or lost.

• Specific Heat– is a property of matter, and different species have

different Specific Heat.

• Specific Heat Capacity is defined as– The heat energy required to raise one gram of a pure

substance one degree Celsius.

Reaction Rates --- 4 Questions

• Students know the rate of reaction is the decrease in concentration of reactants or the increase in concentration of products with time.

•Students know how reaction rates depend on such factors as concentration, temperature, and pressure.

•Students know the role a catalyst plays in increasing the reaction rate .

Reaction Rates• Reaction rates

– how fast a reaction proceeds.

• Some factors will affect reaction rate:

• Temperature of reactants: higher = faster

• Concentration of reactants: greater = faster

• Surface area of reactants: greater = faster– (powders react faster than chunks)

• Catalyst presence: catalysts make rxns faster– Catalysts reduce activation energy!

Chemical Equilibrium --- 4 Questions

• Students know how to use LeChatelier’s principle to predict the effect of changes in concentration, temperature, and pressure. (3 Questions)

• Students know equilibrium is established when forward and reverse reaction rates are equal.

Organic Chemistry and Biochemistry --- 2 Questions

• Students know large molecules (polymers), such as proteins, nucleic acids, and starch, are formed by repetitive combinations of simple subunits.

•Students know the bonding characteristics of carbon that result in the formation of a large variety of structures ranging from simple hydrocarbons to complex polymers and biological molecules.

•Students know amino acids are the building blocks of proteins.

Polymers

• Polymers are long chains or lattices of molecules.• Each molecular building block of polymer is referred to

as a monomer.• Consider the synthetic fabric polyester. Can you

guess what the monomer is?• A monomer is denoted with a bracket and a number,

telling us how many monomer unit make up the average-sized polymer unit. Consider polyethylene a plastic used in milk jugs, the monomer is above.

Nuclear Processes --- 2 Questions

• Students know protons and neutrons in the nucleus are held together by nuclear forces that overcome the electromagnetic repulsion between the protons.

•. Students know the energy release per gram of material is much larger in nuclear fusion or fission reactions than in chemical reactions. The change in mass (calculated by E=mc2) is small but significant in nuclear reactions.

•Students know some naturally occurring isotopes of elements are radioactive, as are isotopes formed in nuclear reactions.

•Students know the three most common forms of radioactive decay (alpha, beta, and gamma) and know how the nucleus changes in each type of decay.

•Students know alpha, beta, and gamma radiation produce different amounts and kinds of damage in matter and have different penetrations.

Isotopes – Nuclides - Radioactivity

• Nuclides– the nucleus of an isotope

• Place the mass above the charge as seen here.

• Nuclides undergo decay: – transformation into different nuclides– Balanced nuclear reactions– Called “Radioactivity”– 3 primary types of decay…

Images from ChemZone

mass

charge

Alpha Decay• Alpha Decay– a helium nucleus is released.

• Alpha particles:– move very slowly– because of their size, can be

blocked with a few pages of paper or human skin

– cause ionization (damaging!)– are positively charged

Images from ChemZone

mass

charge

Alpha Decay occurs in all elements with atomic number above 83.

Beta Decay• Beta Decay

– An electron is ejected from the nucleus

• Beta particles– move fast– can penetrate thick low-

density materials– but can be blocked with

concrete and metals– are negatively charged

Images from ChemZone

Beta Decay occurs when a nucleus has a high neutron-proton ratio.

Gamma Decay• Gamma Decay– High energy photons (gamma rays) are given off.

• Gamma rays– given off as the “spare change” during other

radioactive decays….– extremely penetrating and powerful. Several

inches of lead is required to slow these particles down to a stop.

– Don’t get included in nuclear equations.

Images from ChemZone

Summary of three basic particle decays

No mass

No charge