Chemistry regents in-depth review

Atomic Concepts

• Atoms can be broken down into smaller subatomic particles: electrons, protons, and neutrons.

Subatomic Particles

• Atom's identity is defined by the number of protons, which never changes for a given element.

• Neutral Atom: number of protons = number of electrons.

Atomic Models

• Helps scientists imagine microscopic events on a macroscopic level.

Rutherford's Model

• Most of the atom's mass is in the center (nucleus), which is positive.

• Protons are in the nucleus.

• Most of the atom is empty space.

• Electrons orbit the nucleus.

Bohr's Model

• Protons are in the positive nucleus.

• Electrons revolve around the nucleus in concentric orbits.

Orbital Model (Modern Model)

• Electron Cloud Model.

• Electrons are located in ORBITALS - the most probable location of an electron.

Rutherford's Gold Foil Experiment

• Alpha particles (positively charged) were bombarded at a gold foil.

• Most particles went straight through, but some bounced back.

Conclusions

1. An atom is mostly empty space.

2. An atom has a positively charged nucleus.

Atomic Number & Mass Number

• Atomic Number:

  • Located on the lower-left corner of the element's box on the Periodic Table.

• Mass Number:

  • Located on the upper-left corner of the element's box on the Periodic Table.

  • Equals the total number of particles in the NUCLEUS (PROTONS + NEUTRONS).

  • #\text{ Neutrons} = \text{Mass Number} - \text{Atomic Number}

Isotopes

• Isotopes: atoms of the same element with different numbers of neutrons.

Average Atomic Mass

• The average of all naturally occurring isotopes of a given element.

Principle Energy Levels

• Energy level indicates how far an electron is from the nucleus.

• The first energy level is closest to the nucleus.

• Electron energy increases as levels increase.

  • First principle energy level: holds only 2 electrons.

  • Second principle energy level: holds only 8 electrons.

  • Third principle energy level: holds only 18 electrons.

  • Fourth principle energy level: holds only 32 electrons.

Electron Configuration & Valence

• Electron Configuration:

  • Located on the lower-left corner, below the atomic number.

  • Indicates the number of electrons in each principle energy level.

• Valence Electron:

  • The number of electrons in the last principle energy level.

  • Octet rule: no more than 8 valence electrons.

  • Non-valence electrons: all electrons other than valence electrons.

• Example:

  • Mg 2-8-2: 2 valence electrons, 10 non-valence electrons.

Ground & Excited States

• Ground state: electrons fill the atom in the order 2-8-18-32, as shown on the periodic table.

• Excited state:

  • Electron goes from ground to excited - energy is absorbed.

  • Electron goes from excited to ground - energy is released as the bright-line spectrum (color).

  • Every element releases/absorbs a different amount of energy.

Ions

• Ions: charged particles.

  • Anion: negatively charged ions (gains electrons).

  • Cation: positively charged ions (loses electrons).

Lewis Electron Dot Diagrams

• Show the number of valence electrons in the last principle energy level of an atom.

Nuclear Chemistry

• Ratio of neutrons to protons determines nuclear stability.

• For elements with atomic numbers 1-20, a neutron-to-proton ratio of about 1 results in a stable nucleus.

• Isotopes have the same number of protons but different numbers of neutrons (same atomic number, different atomic mass).

Radioactivity

• Unstable isotope or element decays spontaneously, emitting rays and particles.

Transmutation

• The decay of an atom's nucleus, changing one element into another.

• Mass and charge must be equal on both sides of the equation.

• Table N: decay modes for radioactive isotopes undergoing Natural Transmutation (spontaneous decay).

Natural Transmutation

• Spontaneous decay of the nucleus.

Artificial Transmutation

• Elements made radioactive by bombarding their nuclei with high-energy particles.

• The element will change into another element when the nucleus decays, it doesn't occur spontaneously.

• The atomic # and the mass # have to equal the same thing on both sides.

Separating Alpha, Beta, & Gamma Particles

• Separation via electric or magnetic field.

  • Alpha Particle:

    • Positively charged (2 protons).

    • Deflected toward the negative electrode in an electric field.

  • Beta Particle:

    • Negatively charged.

    • Deflected toward the positive electrode in an electric field.

  • Gamma Rays:

    • No charge.

    • Not deflected.

Nuclear Energy

• Mass is converted into energy during a nuclear reaction.

• Types: fission and fusion.

Fission

• Artificial transmutation.

• A neutron bombards an atom, causing it to split into two or more pieces, releasing energy.

• Controlled chain reaction: nuclear power plants.

• Uncontrolled chain reaction: atomic bombs.

Fusion

• Two nuclei unite to form a heavier nucleus.

• Requires high temperature and pressure (occurs on the sun).

• Releases more energy than fission.

Half-Life

• The time it takes for a sample to decay by half.

• Each radioactive isotope has its own rate of decay.

• Table N: lists half-life rates for common isotopes.

Half-Life Calculations

1. What is the half-life?

   • \frac{\text{Total time elapsed}}{\text{# half-life series}}

2. How much total time elapsed?

   • \text{Half-life} \times \text{# half-life series}

3. How many grams will remain after …?

   • Determine the number of half-life series and successively half the original amount.

4. What fraction remains?

   • (1/2)^{\text{# half-lives}}

     • Knowing the fraction remaining gives you the number of half-life series.

Chemical vs. Nuclear Energy

• Nuclear reactions release more energy than chemical reactions.

Benefits of Nuclear Energy

• Provides a lot of energy (produces the energy of the sun).

• Less $CO_2$ is produced compared to fossil fuels.

Risks of Nuclear Energy

• Radioactive wastes must be stored for >100,000 years without leakage (long half-lives).

• Accidents can cause radioactive spills (mutation/death).

Radioactive Isotopes (Radioisotopes)

• Tracers: Date things.

  • Carbon-14: Living things.

  • Uranium-238 & Lead 206: Non-living things.

• Medical:

  • Isotopes with short half-lives are eliminated by the body quickly.

    • Technetium-99: brain tumors.

    • Iodine-131: thyroid disorders.

    • Radium & Cobalt-60: treatment of cancer.

Risks of Radioisotopes

• Biological Damage: exposure can damage or destroy cells – mutation.

• Long-term storage: must be stored in special containers for a long period of time, is it safe?

• Accidents: 1986 – Chernobyl, 1979—3 Mile Island & Pollution: radioactive materials in the air, water, food, and soil.

Physical Behavior of Matter/Energy

Law of Conservation of Energy

• Energy cannot be created or destroyed, only transferred.

• Forms of energy: heat, chemical, electrical, mechanical, nuclear, potential, kinetic.

Matter

• Anything that has mass and takes up space.

  • Substances: always homogeneous (pure).

  • Mixtures: homogeneous or heterogeneous.

    • Elements: cannot be broken down into simpler substances.

    • Diatomic molecules: elements ($H<em>2$, $O</em>2$, $N<em>2$, $F</em>2$, $Br<em>2$, $Cl</em>2$, $I_2$).

    • Compounds: can be broken down into simpler substances.

    • Binary compounds: composed of 2 elements.

Ways to separate Mixtures

1. Magnet: Separate based on magnetic properties (e.g., sulfur and iron).

2. Distillation: Separate by boiling point (e.g., salt water).

3. Filtration: Use a filter; liquid passes through, solids remain.

States of Matter

• Solids, liquids, gases.

Physical vs. Chemical Changes

• Energy is absorbed or released in a chemical reaction.

• Reactions involving heat energy:

  • Exothermic: energy is released. A + B → C + Heat

  • Endothermic: energy is absorbed. A + Heat → B + C

Phase Changes

• Solid → Liquid → Gas: ENDOTHERMIC.

  • Solid → Liquid = melting.

  • Liquid → Gas = evaporation.

  • Solid → Gas = sublimation (CO2 and I2).

• Gas → Liquid → Solid: EXOTHERMIC.

  • Gas → Liquid = condensation.

  • Liquid → Solid = freezing (solidification).

  • Gas → Solid = Deposition (CO2 and I2).

Entropy

• A measure of the disorder of a system.

• More disorder = more entropy.

• Gases > liquids in disorder.

Measuring Energy

• Energy is measured in joules. 1 kilojoule = 1000 joules.

• Heat energy formula (Table T):

$q = mc\Delta T$

• q: heat.

• m: mass.

• c: specific heat capacity.

• $\Delta T$: change in temperature.

  • If q is negative: exothermic reaction (release heat).

  • If q is positive: endothermic reaction (absorb heat).

  • Heat of vaporization plateau > heat of fusion plateau due to the energy needed during boiling and condensation.

Temperature

• A measure of the average kinetic energy of molecules.

• Higher the temperature = more kinetic energy.

• Heat flows from higher to lower temperature areas until temperatures are the same.

• 0 Kelvin = Absolute 0

• When KE changes PE remains constant and vice versa.

• Know what happens to KE and PE at each interval.

Gases/Gas Laws

• STP = Standard Temperature & Pressure (Table A).

Boiling Point

• Water boils when vapor pressure = atmospheric pressure (Table H).

• Water boils when vapor pressure equals atmospheric pressure (Table H: when water is at 100ºC, atmospheric pressure is 101.3kPa).

• Must know which has the strongest and weakest IMF.

Combined Gas Law

• (TABLE T): An equation that can be used to determine for pressure, volume, or temperature. Nothing remains constant.

$\frac{P<em>1V</em>1}{T<em>1} = \frac{P</em>2V<em>2}{T</em>2}$

• Remember that by stating STP, numerical values are being given!!!!!

• Temperature MUST be in Kelvin

Boyle’s Law

• At constant temperature, the volume of a gas is inversely proportional to pressure.

• $P<em>1V</em>1 = P<em>2V</em>2$

• Doubling the pressure will half the volume.

• 2 times as much pressure = ½ volume of gas

• 3 times as much pressure = 1/3 volume of gas

Charles’s Law

• At constant pressure, volume is directly proportional to Kelvin (absolute) temperature.

• $\frac{V<em>1}{T</em>1} = \frac{V<em>2}{T</em>2}$

• Temperature must be in degrees Kelvin, and therefore, conversion from Celsius to Kelvin may have to occur.

• BOTH BOYLE’S & CHARLES’S LAW CAN BE SOLVED BY USING THE COMBINED GAS LAW and just removing either pressure or temperature from the equation**

Kinetic Molecular Theory

• A model that tells how gases should behave; also called “ideal gas laws”.

• There are 4 key points:

  1. Gas particles are in continuous, random, straight-line motion.

  2. Energy is transferred between colliding particles; total energy remains constant.

  3. Gas particle volume is negligible compared to the space they occupy.

  4. Gas particles have no attraction for each other.

Real Gases

• Deviate from the “ideal gas laws”.

  • Point 3 Deviation: The volume of gas particles is significant. Gas particles do have some volume.

  • Point 4 Deviation: Gas particles do have a force of attraction.

Ideal Gases

• Low pressure & high temperature (particles far apart from each other, and moving very fast)

Real Gases

• High pressure & low temperature (particles close together, and moving slowly by one another)

• Hydrogen and Helium ($H_2$ and $He$) are two REAL gases that act most like IDEAL gases.

• Avogadro’s Hypothesis or Law: Equal volumes of all gases under the same conditions of temperature and pressure have equal numbers of molecules.

• REMEMBER: $6.02x10^{23}$ particles

Periodic Table

Location and Arrangement of Elements

• Periods: Horizontal rows (elements have the same # of PEL’s).

• Groups: Vertical columns (elements have the same # of valence electrons).

• Periodic Law: The properties of the elements are a periodic function of their atomic number

• Elements in the same groups have more similar chemical properties than elements in the same period because they have the same number of valence electrons.

Representative Groups

• Groups 1, 2, and 13-18.

Transition Metals

• Groups 3-12.

• Do not follow the same types of trends as representative groups.

Specific Groups

• Group 1 – Alkali Metals.

• Group 2 – Alkaline Earth Metals.

• Groups 3-12 – Transition Metals: transition metals that have more than one oxidation number form ions that are colored in solution.

• Group 17 – Halogen Group.

• Group 18 – Noble Gases.

Classifying Elements

• Metal: Left of the zig-zag line on the periodic table.

• Non-metal: Right of the zig-zag line on the periodic table.

• Metalloid: On the zig-zag line (except Al and Po).

States of Element at Room Temperature

• Most elements are solid with the following exceptions

  • Gases: 11, 5 are reactive and 6 are non-reactive (reactive gases are $H<em>2$, $O</em>2$, $N<em>2$, $F</em>2$, $Cl_2$; the non-reactive are the Noble Gases).

  • Liquids: Bromine is the only non-metal liquid; Mercury is the only metal liquid.

Reactivity

• Francium is the most reactive metal.

• Fluorine is the most reactive non-metal.

Allotropes

• Different forms of the same element with different molecular or crystalline structures.

• Oxygen: $O<em>2$ and $O</em>3$.

• Carbon can differ in arrangement of atoms:

  • Diamond: every carbon bonded to 4 other carbons = very hard.

  • Graphite: arranged in sheets or layers = “lead” pencils.

  • Coal: no definite pattern.

  • Buckminsterfullerene: rings of 5 and 6 carbon atoms, looks like the outside of a soccer ball (60-70 carbons).

Properties of Metals

• Luster, malleable, good conductors of heat and electricity, are very soluble (metals in group 1 are more soluble than metals in groups 2 and so on).

• Transition metals are harder than metals in group 1 or 2.

• Elements in group 1 are more reactive than elements in group 2.

Properties of Non-metals

• Brittle, lack luster, poor conductors of heat and electricity

• Non-metals are usually gases, molecular solids, or network solids.

Periodic Trends

• Ionization Energy:

  • Energy needed to remove an electron.

    • Ionization energy decreases as you go down a group: Because there are more PEL’s, so the farther away the atom is, the less attraction; thus, it requires less energy to lose electrons.

    • Ionization energy increases as you go across a period: Because there are more protons and more valence electrons as you go across a period and therefore more attraction so requires more energy to lose electrons.

• Electronegativity:

  • The attraction for electrons.

    • Electronegativity decreases as you go down any group: Because there are more PEL’s the farther away you go, the less attraction there is so requires less energy to lose electrons.

    • Electronegativity increases as you go across a period: Because ther are more valence electrons as you go across a period thus there is more attraction so requires more energy to lose electrons.

• Atomic Radius:

  • Is the distance from the nucleus to the outer valence electrons.

    • Atomic radius increases as you go down a group: Because you are adding energy levels.

    • Atomic radius decreases as you go across a period: Because you are increasing the number of protons, there is more attraction as you go across a period and so the distance from the nucleus to the outer PEL decreases slightly

• Ionic Radius:

  • The distance of the nucleus to the outermost valence electron in an ion.

Key Points

• GAIN = GREATER.

• LOSE = LESS.

Nomenclature/Balancing/Reactions

• Reactants → Products

Balancing Equations

• The elements on one side of the equation must equal the elements on the other side of the equation.

Molecular vs. Empirical Formulas

• Molecular formulas: indicates the total number of atoms of each element needed to form a molecule. Ex: $C<em>2H</em>6$ (2 carbons and 6 hydrogens).

• Empirical formula: is the simplest ratio in which atoms combine to form a compound. If the formula is $C<em>2H</em>6$ divide the C and the H by the largest number that all of the elements can be divided by, in this case 2, and the empirical formula would be $CH<em>3$. Try $C</em>6H<em>{12}O</em>6$.

Naming Type 1 Compounds: Ionic

• May occur between the following:

  • Metal (+) & Non-metal (-)

  • Metal (+) & Polyatomic ion (-)

  • Polyatomic ion (+) & Non-metal (-)

  • Polyatomic ion (+) & Polyatomic ion (-)

• NEVER change the name of the polyatomic ions

• The metal keeps its same name

• The non-metal you drop the last few letters and add –ide.

• Criss-cross method: if the oxidation numbers of the elements don’t equal zero, then you must cross the number of the oxidation number, not the charge, in order for the charge to equal 0.

• If you can reduce the subscripts then reduce.

• EX:

  • $CaCl_2 = \text{ calcium chloride}$,

  • $NH_4Cl = \text{ammonium chloride}$,

  • $Li<em>2(SO</em>4) = \text{lithium sulfate}$,

  • $(NH<em>4)</em>2(SO_4) = \text{ammonium sulfate}$

Naming Type 2 Compounds: (Stock System) IONIC

• USE THIS ONLY WHEN THE METAL HAS MORE THAT 1 OXIDATION #

• THE ROMAN NUMERAL INDICATES WHICH OXIDATION # TO USE;

  • Example: I = +1, III = +3, VI = +6

• EX:

  • $Fe<em>2(O)</em>3 = \text{iron (III) oxide}$,

  • $Au_2O = \text{gold (I) oxide}$.

Naming Type 3 Compounds: Covalent

• Use this when you name covalent/molecular compounds (2 NM’s)

• Use the prefixes to name the compounds

• First non-metal keeps its same name and only gets a prefix if there are more than 1

• Second non-metal always gets a prefix, and the ending is –ide.

  • Mono - 1 Hexa - 6

  • Di - 2 Hepta - 7

  • Tri - 3 Octa - 8

  • Tetra - 4 Nona - 9

  • Penta - 5 Deca - 10

Classifying Chemical Reactions

• When two or more chemicals are brought together, a chemical change (reaction) is likely to take place.

• Some evidence that a chemical reaction has occurred include the following:

  1. A color change occurs

  2. A solid forms (precipitate) – INSOLUBLE (Table F)

  3. A gas is released (bubbles form)

  4. Heat and/or a flame are produced (exothermic)

  5. Heat is absorbed (endothermic)

Synthesis Reactions

• Two or more elements or simpler compounds unite to form a compound (BARF)

• $2Na + Cl_2 \longrightarrow 2NaCl$

Decomposition Reactions

• A compound is broken down into two or more simpler compounds (BARF)

• $2NaCl \longrightarrow 2 Na + Cl_2$

Single Replacement Reactions

• A free element (an element alone like Fe) replaces an element that is part of a compound.

• Free element must be more reactive than the element it’s replacing (Table J)

• $Fe + CuSO<em>4 \longrightarrow FeSO</em>4 + Cu$

• $F<em>2 + CuI</em>2 \longrightarrow CuF<em>2 + I</em>2$

Double Replacement Reactions

• Two elements replace each other or switch partners.

• In the example given, the Na and Ag replace each other or switch partners.

• Two new compounds, $NaNO_3$ (sodium nitrite) and $AgCl$ (silver chloride) are formed.

• $NaCl + AgNO<em>3 \longrightarrow NaNO</em>3 + AgCl$

Combustion Reactions

• The reaction of a carbon-based compound with oxygen; the products are carbon dioxide and water (first 6 reactions on Table I)

• $C<em>3H</em>8 + 5O<em>2 \longrightarrow 3CO</em>2 + 4H_2O$

Organic Chemistry

• Organic Chemistry: study of carbon and carbon compounds; there are a lot of carbon compounds. The C atoms bond together to form chains or rings.

Common Characteristics:

1. Generally non-polar (won’t dissolve in water)

2. Soluble in non-polar solvents

3. Non-electrolytes.

4. Low melting points

5. Reactions generally slower than inorganic compounds

6. Reactions have higher activations energies (therefore slower reactions)

• Carbon has 4 valence electrons and therefore can form 4 covalent bonds around them (remember single, double or triple)

• Hydrocarbons: contain carbon and hydrogen atoms. Table P &Q: Table P indicates the prefix used based on the # of carbons

Alkane: (-ane)

• Hydrocarbons are made up of only single bonds, and are therefore considered saturated.

Alkene: (-ene)

• Alkenes have one double bond and classified as unsaturated.

Alkyne : (-yne)

• Alkynes have one triple bond and are considered unsaturated.

• Condensed Formulas: Taking an organic compound, and shortening the chemical formula. Same compound, just written differently.

Alkyl Group: (-yl)

• Also hydrocarbons, known as side chains.

• Have one less hydrogen then a corresponding alkane.

SUMMARY ON HOW TO DRAW CARBON COMPOUNDS

1. Look at the prefix (table P), which will tell you the number of carbons.

2. Put the bonds between the carbons (-ane = single bonds; -ene = double bond; and –yne means a triple bond)

3. The number before –ene and –yne tells you where the double or triple bond is after that carbon atom, Ex: 2-pentene (db after 2nd carbon), or 1-butyne (tb after 1st carbon).

4. When needed, if you have an alkyl group like methyl ($CH_3$), the number before tells you which carbon atom to put it on; Ex: 2-methyl (put the methyl on the 2nd carbon).

5. Put all of the appropriate hydrogen atoms around the carbon atoms if they belong there.

Isomers

• Compounds that have the same molecular formula but different structural formula.

• The rule to naming these structures is to count the number of carbon atoms in the longest unbroken chain. You want to use the lowest possible number, so this means you might have to count from the left to right or from right to left.

Other Organic Compounds & Functional Groups (Table R):

Alcohols: (OH)

• The OH in an alcohol is a hydroxyl, and the OH- in a base is a hydroxide ion; ALCOHOLS ARE NOT BASES.

• To name an alcohol, drop the –e from the corresponding alkane and all –ol.

• OH can be placed on any carbon and therefore a number is required to indicate on which carbon it’s on when there are 3 or more carbons in a chain.

• Diol (2 OH are present); triol (3 OH are present).

Ether:

• You name this compound by using the word –ether at the end and use the name of the alkyl groups attached at both ends.

• Alcohols and Ethers make isomers of each other when there are the same number of carbons atoms.

Aldehyde: (CHO)

• Named by dropping the final –e in an alkane and adding –al.

• Found at the end, no number required.

Ketone:

• Ketones are named by dropping the final –e from the corresponding alkane and adding –one.

• The double bonded oxygen MUST be on an inside carbon, and therefore, requires a number to indicate on which carbon it’s attached to on the chain.

• Aldehydes and Ketones make isomers of each other when there are the same number of carbon atoms.

Organic Acids: (COOH)

• Named by dropping the final -e from the corresponding alkane and adding –oic acid.

• Functional group at the end, and therefore, no number is required.

Halides: (F, Cl, Br, I)

• Name it by using a number to state which atom the halogen is being attached to, then use the prefix for that halogen (fluoro-, chloro-, bromo-, iodo-) and end with the appropriate alkane.

• There can be multiple halides so it’s necessary to use a number to indicate which carbon(s) they are on. Use prefixes like di, tri, etc … if there are more than one of the same kind of halogen.

Amine:

• Named by dropping the final –e in the hydrocarbon and adding –amine.

• Can be placed on any carbon, so a number is required.

Amide:

• Named by dropping the final –e in the hydrocarbon and adding –amide.

• Functional group at the end, therefore no number is required.

Ester: (COO)

• Esterfication – Alcohol + Acid → Ester + Water

How to name:

1. Name the part that came from the alcohol first using an alkyl name depending on how many carbon atoms there were.

2. Name the part that came from the acid second by naming it like an alkane depending on the number of carbons, dropping the –ane, and adding the ending –oate.

Effects of Functional Groups on Boiling Point

1. Alcohols and Acids: Have the highest BP because they are held together with hydrogen bonding.

2. All other functional groups: Next highest BP.

3. Hydrocarbons: Weakest BP.

Examples

• Methanol → Methanal → Methane (highest to lowest BP)

Organic Reactions

Substitution

• Occurs in ALKANES only.

• 2 reactants and 2 products.

Addition

• Happen only to ALKENES & ALKYNES. Breaking a double to a single or a triple to a double.

• 2 reactants and 1 product.

Fermentation

• Glucose is broken down into ethanol and carbon dioxide (Anaerobic Respiration).

Esterfication

• Alcohol + Acid → Ester + Water

Saponification

• Reverse of esterification. This process produces SOAP.

Combustion

• First 6 reactions on Table I.

• Organic Compound + $O<em>2 \longrightarrow CO</em>2 + H_2O$

Polymerization

• Involves smaller molecules joining together to form one big molecule. A polymer is a large molecule made up of multiple monomers.

Condensation

• Polymerization by dehydration synthesis (removing water) to form a polymer. Ex: nylons, polyester. Naturally occurring polymers include starches and protein.

Addition

• Polymerization joining together by breaking a double or triple bond.

Chemical Bonding

• Electronegativity: an atom’s attraction for electrons in a bond.

• Metals tend to have lower EN and Non-metals have higher EN.

• IONIC BONDS: formed when a metal transfers one or more electrons to a non-metal to form ions.

  • Opposite charges attract. Ionic bonds are ALWAYS POLAR BONDS.

Ionic Bond Properties

1. Hard

2. Good conductors of electricity – ONLY IN LIQUID OR AQUEOUS PHASE.

3. High melting and boiling points

4. Dissolve in polar substances: “Likes Dissolve in Likes”

• Ionic compounds that have at least one Polyatomic ion in it is said to have BOTH Ionic and Covalent Bonds.

Covalent Bonds

• Formed when two atoms (both non-metals) share electrons.

• Also called molecular bonds.

• PROPERTIES OF COVALENT BONDS

  1. Soft

  2. Poor conductors of heat/electricity

  3. Low melting and boiling points

Coordinate Covalent

• Formed when one atom donates a pair of electrons. $NH<em>4 +$ and $H</em>3O^+$.

Network Solids

• Solids that have covalent bonds between atoms linked in one big network or one big macromolecule with no discrete particles. Examples: Diamond (C), silicon carbide (SiC), and silicon dioxide ($SiO_2$).

• PROPERTIES of network solid substances that have covalent bonds

  1. Hard

  2. Poor conductors of heat and electricity

  3. High melting points

Metallic Bond

• “Mobile Electrons”. Why metals are good conductors in the solid phase.

• INTERMOLECULAR FORCES: forces of attraction between molecules

Hydrogen Bonding

• Strong intermolecular force connects one water molecule with another (also other examples)

• Hydrogen bonds are also formed when hydrogen is covalently bonded to a small, highly electronegative atom such as F, O, or N. Examples: $H<em>2O$, $NH</em>3$, and $HF$.

Dipole Attraction

• Same concept as Hydrogen Bond but for other polar molecules (asymmetrical with partially (+) and (-) ends. Ex: $H<em>2S$, $PCl</em>3$, $HCl$.

Dispersion Forces

• (AKA: van der Waals forces) Weak intermolecular forces between SYMETRICAL NON-POLAR MOLECULES.

• No $\delta +$ or $\delta -$.

• Increases with GFM. As you go down group 17 these forces increase and BP increase. This is why $F<em>2$ and $Cl</em>2$ are gases, $Br<em>2$ is a liquid, and $I</em>2$ is a solid.

Molecule-Ion Attraction

• Attraction between the ions of an ionic compound such as salt and molecules of water (or other polar liquids).

• “Likes Dissolve in Likes”

• Greater EN $\Delta$ = More Polar = Stronger IMF = Higher MP/BP

Moles/Stoichiometry

Moles/Stoichiometry

• A mole is a standard measurement of the amount of substance; it is the GFM (gram formula mass).

Stoichiometry

• The study of quantitative relationships that are derived from balanced chemical equations.

Key Points:

• The coefficients tell you how many moles there are.

• The coefficients are not the number of grams present in the equation.

• You can’t go directly from the equation to the number of grams of each substance.

Moles Formula

• Number of moles = Given grams/GFM

• If using grams, use Table S

• If given moles, set up a proportion using the coefficients in the equation as the new moles.

Math of Formulas

Percent Composition

• Mass % = mass of part/mass of whole x 100

• If they ask for the mass %, look at the GFM (molar mass) of the compound, and look at how many moles there, Ex: finding the % mass of O in $Fe<em>2O</em>3$ = you have 3 moles of O x mass of O x 100/GFM of the whole compound.

Parts Per Million (PPM)

• PPM = grams of solute/grams of solution x 1,000,000

• PPM = mL of solute/mL of solution x 1,000,000

Hydrates

• Compounds that have water within their crystalline structure (waters of hydration)

• $CuSO*4 * 5H*2O$ (cupric sulfate pentahydrate)

• To solve for % of water, solve for each mass and then divide mass of water by the whole hydrate.

Solutions

Properties of Solutions

• Solutions must be homogenous

• Smallest particle size (0-1nm)

• Particles are not visible and don’t settle

• Particles pass through filter paper

• Do not show the Tyndall effect

Molarity

• Concentration = Molarity = moles of solute/liter of solution (MEMORIZE THIS)

Solubility

• How much solute can dissolve in a solvent and depends on-

  • Temperature

  • Pressure

  • Nature of solute and solvent

• SOLUBILITY CURVE = Table G

• Table F = solubility guidelines (memorize)

Solubility Key Points:

• Unsaturated: can dissolve more solute

• Saturated: on the line (holds max amount of solute)

• Supersaturated: holds more solute than it should (under the line)

• Non-polar dissolves non-polar, and polar dissolves polar (“Likes Dissolve in Likes”)

• Gases are more soluble at low temperatures and high pressures

• Solids are more soluble at high temperatures

Dissolving Solids

• Increasing the surface area, temperature, and stirring increases the rate of dissolving

Colligative Properties

• Adding a solute to a solvent will- Lower the freezing point.

  • Raise the boiling point.

  • The greater the concentration, the greater the effect (Adding salt increases the BP in water, the more the salt the higher the BP)

Kinetics & Equilibrium

Collision Theory

• In order for a reaction to occur molecules must collide with

  • Sufficient energy.

  • Proper orientation.

Reaction Rate

• Rate of chemical reaction depends on-

  • Concentration (increase the concentration, the faster the rate [exception = zero order]).

  • Temperature (increase the temperature, the faster the rate).

  • Pressure (more pressure, the faster the rate).

  • Catalysts (the presence of a catalyst will speed up the reaction, without being consumed)

  • NATURE OF REACTANTS IN A CHEMICAL REACTION

• Surface Area (the more surface area there is, the faster the reaction)

Energy of Activation (Ea)

• The amount of energy needed to form the activated complex

Catalyst

• Provide a new pathway with a lower activation energy

• Lowers the activation energy (Ea) to speed up the reaction

Potential Energy Diagrams

• Know which way the graph will go for exothermic and endothermic graphs

• Be able to label heat of reaction (ΔH), activated complex, activation energy, etc

Nature of Reactants Key Points:

• Solids react slower than liquids

• Liquids react slower than gases

• Ionic bonds react faster than covalent bonds

Equilibrium

• Rate of forward reaction = rate of reverse reaction.

• Concentrations are constant BUT NOT EQUAL.

• System is closed; no matter enters or leaves.

Le Chatelier’s Principle

• If a system at equilibrium is subjected to a stress, the equilibrium will shift in the direction that relieves the stress.

Types of Stress

• Concentration: If you add more reactants, the reaction shifts to products to counteract the stress. If you add more products, the reaction shifts to reactants to counteract the stress.

  • Pressure: Only effects GASES. If you increase the pressure, the reaction will shift to the side with less moles or less gas molecules. If you decrease the pressure, the reaction will shift to the side with more moles or more gas molecules.

  • Temperature: Increasing temp will favor endothermic reaction. Decreasing temp will favor exothermic reaction.

Acids, Bases, & Salts

Properties of Acids

• Sour taste

• Electrolytes (conduct electricity – due to the presence of mobile ions).

• Turn litmus paper RED.

• Have a pH less than 7.

• React with active metals to form hydrogen gas.

• Neutralize bases.

Properties of Bases

• Bitter taste

• Electrolytes (conduct electricity – due to the presence of mobile ions).

• Turns litmus paper BLUE.

• Have a pH greater than 7.

• Feels slippery.

• Neutralizes acids.

Acid-Base Theories

Arrhenius Theory

• Acids produce $H^+$ ions as the only positive ion in water. $HCl ento H^+ + Cl^-$

• Bases produce $OH^−$ ions as the only negative ion in water $NaOH ento Na^+ + OH^-$

Bronsted - Lowry Theory

• Acids donate protons ($H^+$).

• Bases accept protons ($H^+$).

• $HCl + H<em>2O ento H</em>3O^+ + Cl^-$

• $HCl$ donates a proton (so it’s an acid) & $H*2O$ accepts a proton (so it’s a base)

Titration

• A process used to determine the unknown concentration of an acid or base by adding a known amount of the opposite (acid or base).

• At the endpoint, the number of moles of $H^+$ ions is equal to the number of moles of $OH^−$ ions. (Neutralization has occurred).

Acid - Base Indicators

• Changes color at different pH values.

• Table M: Common Acid-Base Indicators.

pH Scale

• Measures degree of acidity.

• Range is from 0-14.

• pH of 7 is neutral.

Neutralization

• A reaction between an acid and a base to produce water and a salt. $Acid + Base ento Salt + H*2O$

Salts

• Are ionic compounds consisting of a positive ion from a base and a negative ion from an acid.