NAME:______________________________________________ Period ____

CORE CURRICULUM CHECK-OFF LISTREGENTS CHEMISTRY

The following information is everything you need to know and be able to do to attain Mastery of the Regents Chemistry Curriculum. The bold, underlined words are important vocabulary words that you should be able to define and use properly in explanations. This is a study guide for what you will be tested on throughout the year. The objectives are divided into categories of “Knowledge” (what you have to know) and “Application” (what you have to be able to do). Check off each objective when you fully understand it. If you do not understand an objective, ask questions before you are tested on it.

I. MATHEMATICS ANALYSIS & GRAPHING

1.

o Identify independent and dependent variables in an experiment and correctly plot them on an axis

Example hypothesis: Chemistry students who do their homework will have higher test scores than students who do not do their homework.

X-axis (horizontal): the independent variable is the one that is manipulated by the experimenter. (“The one I change.” – do homework/not do homework)

Y-axis (vertical): the dependent variable is the one that changes based on the independent variable. (The data you collect – test scores)

2.

o Express uncertainty in measurement by properly using significant figureso Identify the number of sig. figs. in datao Round to the correct number of sig. figs. on calculations:

Addition and Subtraction (round to least precise [farthest left] place value)

Multiplication and Division (round to the lowest # of digits in data) Combined Rules (Add/Subtract first, then Multiply/Divide)

3.

o Identify relationships between variables from data tables and graphs (direct or inverse relationships)

4.

o Understand what is meant by conditions of Standard Temperature and Pressure (STP) (Table A)

5.

o Recognize and convert between units on various scales of measurement (Tables C, D, and T)

Temperature: Celsius ↔ Kelvin Length: meters ↔ centimeters ↔ millimeters Mass: grams ↔ kilograms Pressure: kilopascals ↔ atmospheres Thermal Energy: joules ↔ kilojoules Amount of Substance:

GFM = 1 mole = 6.02x1023 particles

o grams ↔ moles ↔ atoms or molecules

6.

o Use the density equation on Table T to solve for density, mass, or volume, given the other two values

7. o Calculate percent error (Table T)

II. ATOMIC CONCEPTSKnowledge Application

1.

o Atoms are the basic unit (building block) of matter.

o Atoms of the same kind are called elements.

o The modern model of the atom has developed over a long period of time through the work of many scientists.

o Explain what happened during the gold-foil experiment and what it showed.

Gold foil was bombarded (hit) with positively charged alpha particles. Most alpha particles passed through the gold foil, but some were deflected. This showed that:

1. the atom is mostly empty space 2. the nucleus is small, positively charged, and located in the center of the atom

2.

o The three subatomic particles that make up an atom are protons, neutrons, and electrons.

o The proton is positively charged, the neutron has no charge, and the electron is negatively charged.(This is also referenced on Table O. Check it out!)

3.

o Each atom has a nucleus with an overall positive charge, made up of protons and neutrons.

o The nucleus is surrounded by negatively charged electrons.

o Determine the nuclear charge of an atom. (equal to the number of protons in the nucleus)

4. o Atoms are electrically neutral, which means that they have no charge (# protons = # electrons)

o Ions are atoms that have either lost or gained electrons and are either positively or negatively charged

o When an atom gains one or more electrons, it becomes a negative ion and its radius increases.

o When an atom loses one or more electrons, it becomes a positive ion

o Determine the number of protons or electrons in an atom or ion when given one of these values. (Periodic Table)

o Compare the atomic radius and ionic radius of any given elementEx: A chloride ion has a larger radius than a chlorine atom because the ion has an extra electron. A sodium ion has a

and its radius decreases.smaller radius than a sodium atom because the ion has lost an electron.

5.

o The mass of each proton and each neutron is approximately equal to one atomic mass unit (AMU).

o An electron is much less massive (has almost no mass) compared to a proton or neutron.

o Calculate the mass of an atom given the number of protons and neutrons

o Calculate the number of neutrons or protons, given the other value

6.o In the modern model of the atom, the WAVE-MECHANICAL MODEL

(electron cloud), the electrons are in orbitals (clouds), which are defined as regions of most probable electron location.

7.o Each electron in an atom has a specific amount of energy.o Electrons closest to the nucleus have the lowest energy. As an electron

moves away from the nucleus, it has higher energy.

8.

o Electron configurations show how many electrons are in each orbital.

o When all of an atom’s electrons are in the orbitals closest to the nucleus, the electrons are in their lowest possible energy states. This is called the ground state.

o When an electron in an atom gains a specific amount of energy, the electron is at a higher energy state (excited state)

o Distinguish between ground state and excited state electron configurations. (Be careful to keep the same number of electrons when writing the excited state.)

II. ATOMIC CONCEPTS (CONTINUED)Knowledge Application

9.

o When an electron returns from a higher energy (excited) state to a lower energy (ground) state, a specific amount of energy is emitted. This emitted energy can be used to identify an element.

o The flame test is an example of the bright-line spectrum visible to the naked eye. The color can determine the identity of a positive ion in a compound.

o Identify an element by comparing its bright-line spectrum to given spectra

10.

o The outermost electrons in an atom are called the valence electrons . In general, the number of valence electrons affects the chemical properties of an element.

o Draw a Lewis electron-dot structure of an atom.

o Distinguish between valence and non-valence electrons, given an electron configuration

11.

o Atoms of an element that contain the same number of protons but a different number of neutrons are

o Calculate the number of neutrons in an isotope of an element given the isotope’s mass

called isotopes of that element.

12.

o The average atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.

o Given an atomic mass, determine the most abundantisotope

o Calculate the atomic mass of an element, given the masses and abundance of naturally occurring isotopes

III. THE PERIODIC TABLEKnowledge Application

1.

o The placement or location of an element on the Periodic Table gives an indication of physical and chemical properties of that element.

o The elements on the Periodic Table are arranged in order of increasing atomic number.

o Explain the placement of an unknown element in the periodic table based on its properties

2.

o The number of protons in an atom (atomic number) identifies the element. This goes at the bottom left corner of the symbol for that element.

o The sum of the protons and neutrons in an atom (mass number) identifies an isotope. The mass number is placed at the top left corner of the symbol for an element OR is placed after the element symbol or name and a dash.Ex: Three different ways to write carbon with a mass number of 14: C, carbon-14, or C-14

o Interpret and write isotopic notations. Ex:C-12, C-13, and C-14 are isotopes of the element carbon

3.o Elements are classified by their properties,

and located on the periodic table as metals, metalloids, nonmetals, or noble gases.

o Identify the properties of metals, metalloids, nonmetals and noble gases

o Classify elements as metals, metalloids, nonmetals, or noble gases by their properties

oIII. THE PERIODIC TABLE (CONTINUED)Knowledge Application

4. o An element’s atomic radius, first ionization energy, and electronegativity determine its physical and chemical properties

5. o Substances can be differentiated by their physical properties.

o Physical properties of substances include melting point, boiling point, density,

o Identify and give examples of physical properties

o Describe the states of

conductivity, malleability, solubility, and hardness.

the elements at STP (solid, liquid, or gas). (Table S)

6.

o Substances can be differentiated by chemical properties.

o Chemical properties describe how an element behaves during a chemical reaction and include reactivity, flammability, and toxicity.

o Identify and give examples of chemical properties

o Describe the difference between physical and chemical properties of substances

7.

o Some elements exist as two or more forms in the same phase. These forms differ in their molecular or crystal structure and therefore in their properties. The word to describe this phenomenon is ALLOTROPE.

o Ozone and oxygen gases are allotropes of each other. Ozone is O3 and it is very dangerous to our health. Oxygen gas is O2 and we need it to survive.

o Diamonds and graphite (better known as pencil lead) are both forms of the element carbon. They have different molecular structures and very different properties.

8.

o For Groups (also called families) 1, 2, and 13-18 on the Periodic Table, elements within the same group have the same number of valence electrons (helium is the exception) and therefore similar chemical properties.

o Elements in the same Period (row) have the same number of principal energy levels (shells) which contain electrons.

o Determine the group of an element, given the chemical formula of a compound Ex: A compound has the formula XCl2, element X is in Group 2

o Determine the number of energy levels containing electrons given an element’s Period and vice versa

9.

o The succession of elements within the same GROUP (top to bottom) demonstrates characteristic trends: differences in atomic radius, ionic radius, electronegativity, first ionization energy, and metallic/nonmetallic properties.

o Compare and contrast properties of elements within a group or a period for groups 1, 2, and 13-18 on the periodic table

10.

o The succession of elements across the same PERIOD (left to right) demonstrates characteristic trends: differences in atomic radius, ionic radius, electronegativity, first ionization energy, and metallic/nonmetallic properties.

o Understand and be able to explain the trends in terms of nuclear charge and electron shielding

IV. MATTER, ENERGY, & CHANGEKnowledge Application

1.

o Matter is anything that has mass and volume (takes up space).

o Matter cannot be created nor destroyed, only transformed.

o Matter is classified as a pure substance (element or compound) or as a mixture of substances.

o Identify specific examples of matter as an element, compound, or mixture

IV. MATTER, ENERGY, & CHANGE (CONTINUED)Knowledge Application

2.

o Energy is not matter (does not have volume)

o Energy can exist in different forms, such as kinetic, potential, thermal (heat), sound, chemical, electrical, and electromagnetic.

o Energy cannot be created or destroyed, only transformed.

o Distinguish between matter and energy

3.

o During a physical change, particles of matter are rearranged. Examples of physical changes include freezing, melting, boiling, condensing, dissolving, crystallizing, and crushing into a powder

o During a chemical change, NEW substances are formed with new properties. Examples of chemical changes include combustion (burning), rusting, and neutralizing an acid or base.

o Energy can be absorbed or released during physical and chemical changes.

o Differentiate between physical and chemical changes in matter

o Identify and give examples of physical changes and chemical changes in matter

4.

o A pure substance (element or compound) has a uniform composition and constant properties throughout a given sample, and from sample to sample.

o Mixtures are composed of two or more different substances that can be separated by physical means.

o Draw and interpret particle diagrams for elements, compounds, and mixtures

5.

o Elements are substances that are composed of atoms that have the same atomic number. Elements cannot be broken down by chemical change.

6.

o A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion.

o A compound can be broken down by chemical means, such as during a chemical reaction.

o Two major categories of compounds are ionic and molecular (covalent)

o Describe differences in ionic and molecular/covalent compounds

o Identify a compound as ionic or molecular/covalent compound given its properties

compounds.o A chemical compound can be represented

by a specific chemical formula and assigned a name based on the IUPAC system.

o Name compounds based on their chemical formulas

o Determine the formula of a compound given its name

7.

o When different substances (elements or compounds) are mixed together and do NOT chemically react, a mixture is formed.

o The amounts of substances in a mixture can vary. Each substance in a mixture retains its original properties.

o The composition of a mixture can vary. If the substances are uniformly (evenly) distributed throughout the mixture, it is called a homogenous mixture. If the substances are unevenly distributed, it is called a heterogeneous mixture.

o Interpret particle diagrams as showing homogeneous or heterogeneous mixtures

o Give examples of homogeneous and heterogeneous mixtures

8.

o Differences in properties such as density, particle size, molecular polarity, boiling point, freezing point, and solubility allow physical separation of the components of the mixture.

o Describe the processes of filtration, distillation, and chromatography and the types of mixtures they are used to separate

V. CHEMICAL BONDINGKnowledge Application

1.

o Atoms bond with other atoms to gain a stable electron configuration.

o Noble gases are already stable and tend to not bond/react.

o Determine the noble gas configuration an atom will achieve when bonding

2.

o When a chemical reaction takes place, existing bonds must be broken in order for new bonds (and new compounds) to be formed.

o When a bond is broken, energy is absorbed. When a bond is formed, energy is released.

3.

o Electron-dot diagrams (Lewis structures) are used to represent the valence electron arrangement in elements, ions, and compounds.

o Draw Lewis dot structures for any given element, ion, or compound

4.

o Chemical bonds are formed when valence electrons are: transferred from one atom to another (ionic) shared between atoms (covalent)

o Identify the type of bonding in a compound (ionic or covalent), given the elements that make it up

o Demonstrate bonding concepts using Lewis Dot Structures (electron-dot diagrams) for ionic and covalent compounds

Ionic compounds – after the transfer of electrons, the positive ion should have no dots.

mobile within a metal (metallic)

o Metals tend to react with nonmetals to form ionic bonds.

o Nonmetals tend to react with other nonmetals to form molecular (covalent) bonds.

o Ionic compounds containing polyatomic ions have both ionic AND covalent bonds.

The negative ion should have 8 dots around it. Put brackets around the ions. Check the periodic table to find the charge associated with the ion and place this charge outside of the brackets. *Be sure to include coefficients if there are more than one of the same kind of ion.*Covalent compounds – each atom in a covalent compound must end up with 8 dots around it – except for hydrogen (only 2 dots). *If you use lines, remember that one line represents 2 electrons being shared.*

5.

o In a multiple covalent bond, more than one pair of electrons is shared between two atoms.

o Draw electron-dot diagrams and give examples of molecules with multiple covalent bonds

6.

o Electronegativity indicates how strongly an atom of an element attracts electrons in a bond.

o The electronegativity difference between two bonded determines the degree of polarity in the bond.

o Distinguish between and give examples of nonpolar covalent bonds and polar covalent bondsIf two atoms of the same element share electrons, the bond is nonpolar (ex: H–H)If atoms of two different elements share electrons, the bond is polar (ex: H–Cl)

7.

o Molecular polarity is determined by the distribution of electrons in a covalent compound.

o Symmetrical distribution of electrons results in (nonpolar) molecules (Ex: CO2, CH4 and all diatomic elements)

o Asymmetrical distribution of electrons results in (polar) molecules (Ex: HCl, NH3, H2O)

o Distinguish between bond polarity and molecular polarity

o Draw Lewis Dot Structures for all of the compounds listed to the left, including the diatomic elements

o Determine whether a molecule is polar or nonpolar, given its structure

*SNAP*

VI. CHEMICAL FORMULAS, REACTIONS & STOICHIOMETRYKnowledge Application

1.

o Chemical formulas are used to represent compounds.

o The main types of chemical formulas include: empirical, molecular, and structural.

o Determine the empirical formula from a molecular formula

o An empirical formula is the simplest whole-number ratio of atoms in a compound.

o Molecular formulas are chemical formulas that show the actual ratio of atoms in a molecule of that compound.

o Structural formulas can also be used to represent covalent compounds. These use lines to show covalent bonds between atoms and also show the geometrical arrangement of atoms in the compound.

o Draw structural formulas for covalent (molecular) compounds

2.

o One mole of any substance is equal to 6.02 x 1023 pieces of that substance.

o The formula mass of a compound is equal to the sum of the atomic masses of its atoms (units are atomic mass units)

o The molar mass (gram-formula mass) of a substance is equal to the formula mass in grams – hence “gram-formula mass”.

o The mass of one mole of any substance is equal to its molar mass (gram-formula mass).

o Calculate the molar mass (gram-formula mass) of a substance

o Determine the molecular formula, given the empirical formula and the molar mass

o Determine the number of moles of a substance, given its mass and vice versa

3.

o The percent composition by mass of each element in a compound can be calculated mathematically.

o Calculate the percent composition of any element in a given compound

o Calculate the percent composition of water in a given hydrate

4.

o Balanced chemical equations show conservation of matter, energy, and charge.

o The coefficients in a balanced equation can be used to determine mole ratios in the reaction.

o Balance equations, given the formulas for reactants and products

o Calculate simple mole-mole ratios, given balanced equations

5.

o Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement.

o Identify the different types of chemical reactions, given their chemical equations

VII. PHYSICAL BEHAVIOR OF MATTERKnowledge Application

1. o Physical properties of substances can be explained in terms of chemical bonds and intermolecular forces.

o Intermolecular forces created by an

o Predict relative melting and boiling points of compounds given information about its

unequal distribution of charge result in varying degrees of attraction between molecules. Hydrogen bonding is an example of a strong intermolecular force that occurs in compounds containing H bonded to F, O, or N atoms.

o Physical properties include malleability, solubility, hardness, melting/freezing point, and boiling point.

chemical bonds or strength of intermolecular forces

o Predict relative strength of intermolecular forces of a compound given its melting and boiling points

VII. PHYSICAL BEHAVIOR OF MATTER – SOLIDS, LIQUIDS, AND GASESKnowledge Application

2.

o The three phases of matter (solids, liquids, and gases) have different properties.

o The structure and arrangement of particles and their interactions determine the physical state (s, l, or g) of a substance at a given temperature and pressure.

o Draw and interpret a particle diagram to differentiate among solids, liquids, and gases

o Explain phase change in terms of the changes in energy and intermolecular distances

3.

o Heat is a transfer of energy (thermal energy) from a body of higher temperature to a body of lower temperature. Heat (thermal energy) is associated with the random motion of atoms and molecules.(Heat flows from HOT materials to COLD materials)

o Temperature is a measure of the average kinetic energy of the particles in a sample of matter. Temperature is NOT energy – it is a measure of heat energy.

o Distinguish between heat energy and temperature in terms of molecular motion and amount of matter

o Convert between degrees Celsius and degrees Kelvin (Table T)

4.

o The concepts of kinetic and potential energy can be used to explain physical processes that include: fusion (melting), solidification (freezing), vaporization (boiling, evaporation), condensation, sublimation, and deposition.

o The kinetic and potential energy changes involved in these physical processes can be illustrated in a heating curve or cooling curve.

o Identify areas of heating and cooling curves that show changes in kinetic and potential energy, heat of vaporization, heat of fusion, and phase changes

o Calculate the heat involved in a phase or temperature change of a sample of matter using Tables B & T and/or a given heating or cooling curve

5. o Physical processes like phase changes can be exothermic or endothermic.

o Distinguish between endothermic and exothermic phase changes

6. o Entropy is a measure of the randomness or o Compare the entropy of

disorder of a system. A system with greater disorder has more entropy. different phases of matter

7.

o The concept of an ideal gas is a model to explain behavior of gases.

o A real gas is most like an ideal gas when the real gas is at low pressure and high temperature.

o Given a choice of pressure and temperature conditions, identify those under which gases behave most ideally and/or least ideally

8.

o The Kinetic Molecular Theory (KMT) states that, for an IDEAL gas, all gas particles

a. are in random, constant, straight-line motionb. are separated by great distances relative to their size (have negligible

volume)c. have no attractive forces between themd. have collisions that may result in a transfer of energy between them,

but the total energy of the system remains constant

9.

o The Kinetic Molecular Theory (KMT) describes the relationships of pressure, volume, temperature, velocity, frequency, and force of collisions among gas molecules.

o Explain the gas laws in terms of KMT.

o Solve problems using the combined gas law (Table T)

o Recognize and draw graphs showing P vs. T, V vs. T, and P vs. V

10.

o Equal volumes of gases at the same temperature and pressure contain an equal number of particles.

VII. PHYSICAL BEHAVIOR OF MATTER – AQUEOUS SOLUTIONSKnowledge Application

11.

o Physical processes, such as a compound dissolving in a solution, can be exothermic or endothermic.

o Interpret ∆H values for physical processes given in Table I

12. o A solution is a homogeneous mixture of a

solute dissolved in a solvent.

o Identify the solute and the solvent in a given solution

o Give examples of different types of solutions

o The solubility of a solute in a given amount of solvent is dependent on the temperature, the pressure, and the chemical natures of the solute and solvent.

o General rules: a. solubility of a solid increases as

temperature increases (direct relationship)

b. solubility of a gas decreases as temperature increases (inverse relationship)

c. solubility of a gas increases as pressure

o Predict the effect of temperature, pressure, and nature of solvent on solubility for a given solute

o Use a solubility curve to distinguish among unsaturated, saturated, and supersaturated solutions

o Calculate the amount of a specific solute dissolved at

increases (direct)d. “like dissolves like” – polar solvents

dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes

different temperatures using Table G

13.

o Many chemical reactions happen in solution. When different ionic compounds are mixed together in the same solution, a double replacement reaction may occur and a stable precipitate (insoluble/solid compound) may form.

o Use Table F (Solubility Guidelines) to determine a compound’s solubility

o Determine if a precipitate will form when ionic compounds are mixed in solution

o Write and balance chemical equations for double replacement reactions

14.

o The concentration of a solution may be expressed as: molarity (M), percent by volume (%v/v), percent by mass (%m/v), or parts per million (ppm).

o Calculate solution concentrations in molarity (M), percent by volume, percent by mass, or parts per million (ppm)

o Describe how you would prepare a solution from scratch, given the desired molarity

o Describe how you would dilute a solution of known concentration (must use the equation M1V1 = M2V2)

15.

o The addition of a nonvolatile solute to a solvent causes the boiling point of the solution to increase and the freezing point of the solution to decrease.

o The greater the concentration of solute particles, the greater the increase in b.p. and decrease in f.p.

o Compare the freezing and boiling points of solutions of different concentration

VIII. KINETICS AND EQUILIBRIUMKnowledge Application

1.

o The Collision Theory states that a chemical reaction is most likely to occur if reactant particles collide with the proper energy and orientation.

2.

o The rate (speed) of a chemical reaction depends on several factors: temperature, concentration, nature of reactants, surface area, and the

o Use the Collision Theory to explain how factors such as temperature, surface area, and concentration influence the rate of

presence of a catalyst.o Ionic compounds generally react

faster than covalent (molecular) compounds

o A catalyst provides an alternate reaction pathway, which has lower activation energy than an uncatalyzed reaction.

reactionEx: Increasing the temperature, surface area, or concentration all lead to an increase in the rate of a reaction because they all increase the number of effective collisions between reactant particles.

o Explain, in terms of the number of bonds broken, why ionic compounds generally react faster than covalent compounds

o Explain how a catalyst speeds up a reaction

3.

o Energy released or absorbed during a chemical reaction can be represented by a potential energy diagram.

o The difference in PE of the products and reactants is called the heat of reaction ( H) H = PE products – PE reactants

o H values for many chemical reactions are listed in Table I

o Read and interpret a potential energy diagram

o Draw and label the following parts of a potential energy diagram for both an endothermic and exothermic reaction PE of reactants and PE of products heat of reaction (H) activation energy (for both the forward and reverse reactions) activation energy with a catalyst present

4.

o Chemical and physical changes can reach equilibrium

o Saturated solutions are examples of systems in physical equilibria (aq ↔ s)

o Distinguish between examples of physical equilibria and chemical equilibria

5.

o At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction and the measurable quantities of reactants and products remain constant at equilibrium*CARE*

o Describe what is happening to the concentrations or amounts of reactants and products in a system at equilibrium

o Describe the rates of opposing reactions in a system at equilibrium

6.

o LeChatelier’s principle can be used to predict the effect of a stress (such as a change in pressure, volume, concentration, or temperature) on a system at equilibrium.

o According to LeChatelier’s principle, a system at equilibrium will “shift”

o Describe, in terms of LeChatelier’s principle, the effects of stress on a given system at equilibrium, including: Changing the

temperature/heating/cooling Changing the concentration of

a reactant or product

to reduce the effects of a stress placed on the system. It will “shift” AWAY from an INCREASE and will “shift”toward a decrease in concentration or temperature (“shift” means that either the forward or the reverse reaction will be “favored” (go faster) until the rates are again equal and equilibrium is re-established).

o Changing the pressure or volume only affects systems that contain gases

Changing the pressure or volume (this affects systems involving gases)

o Also be able to explain why any shifting occurs in terms of Collision Theory

7.

o Systems in nature tend to undergo changes toward lower energy and higher entropy.

IX. ORGANIC CHEMISTRYKnowledge Application

1.

o Organic compounds contain carbon atoms which bond to one another in chains, rings, and networks to form a variety of structures.

o Organic compounds are named using specific IUPAC rules and can be represented using molecular formulas, structural formulas, or condensed structural formulas.

2.

o Hydrocarbons are organic compounds that contain only carbon and hydrogen. *The C–H bonds are considered to be nonpolar covalent bonds.*

o Saturated hydrocarbons contain only single C–C bonds.

o Unsaturated hydrocarbons contain at least one double or triple carbon-carbon bond.

o Draw structural formulas for alkanes, alkenes, alkynes, given their IUPAC names

o Name a hydrocarbon (IUPAC rules), given its molecular formula, structural formula, or condensed structural formula

o Identify whether a hydrocarbon is saturated or unsaturated, given its IUPAC name, structural formula, condensed structural formula, or general formula (see Table Q)

3.

o In a multiple covalent bond, more than one pair of electrons is shared between two atoms. In a structural formula, each line represents TWO shared electrons

o Determine the TOTAL number of electrons shared in a covalent bond

o Determine the number of electron PAIRS (lines) shared in a covalent bond

4.

o A functional group is a group of atoms attached to an organic compound that gives distinct physical and chemical properties to organic compounds having that group attached to it.

o Organic acids, alcohols, esters,

o Identify different kinds of functional groups

o Classify an organic compound based on its structural formula, condensed structural formula, or IUPAC name

o Draw a structural formula with the

aldehydes, ketones, ethers, halides, amines, amides, and amino acids are types of organic compounds that differ in the type of functional group they have.

o Compounds that have the same functional group have similar physical and chemical propertiesEx: all esters have pleasant odors, all organic acids donate H+ ions in solution, all alcohols have low boiling points, etc.

functional group(s) on a straight chain hydrocarbon backbone, given the correct IUPAC name for the compound

o Name any of these organic compounds, given their structural or condensed structural formulas (see Table R)

5.

o Isomers are organic compounds that have the same molecular formula, but different structures and properties.

o Determine if two compounds are isomers, given their molecular formulas, structural formulas, condensed structural formulas, or names

6.

o Types of organic reactions include: polymerization, substitution, fermentation, addition, combustion, esterification, and saponification. *P.S. FACES*

o Identify types of organic reactions, given balanced chemical equations

o Determine missing reactants or products in a balanced equation, given the type of reaction.

X. OXIDATION-REDUCTION REACTIONSKnowledge Application

1.

o An oxidation-reduction (redox) reaction involves the transfer of electrons from one species (element or ion) to another.

o The number of electrons lost equals the number of electrons gained (conservation of charge)

o Determine the number of moles of electrons lost or gained in a redox reaction, given the other value

2.

o Oxidation numbers (states) can be assigned to atoms and ions.

o Changes in oxidation numbers indicate that a redox reaction has occurred.

o Assign oxidation states to atoms and ions

o Determine if a reaction is a redox reaction given the reaction equation (Hint: any reaction in which an element is alone (uncombined with another element) on one side, but in a compound on the other side – it’s a redox reaction!)

3.

o Losing Electrons is Oxidation (LEO)o Gaining Electrons is Reduction

(GER)o Oxidized and reduced species are

ALWAYS on the LEFT (reactants) side of the equation.

o Determine which species undergoes reduction (oxidation state goes down) and which species undergoes oxidation (oxidation state goes up)

4.

o An oxidation half-reaction shows which species is oxidized and the number of electrons it loses (electrons go on the right side of the arrow)

o A reduction half-reaction shows which species is reduced and the number of electrons it gains (electrons go on the left side of the arrow)

o Determine if a given half-reaction is showing oxidation or reduction

o Write and balance oxidation and reduction half-reactions (*remember conservation of mass and charge – multiply one or both of the half-reactions to make electrons lost = electrons gained if they are not equal at first)

5.

o An electrochemical cell can either be a voltaic cell (a battery) or an electrolytic cell.

o In both voltaic and electrolytic cells oxidation occurs at the anode (An Ox) reduction occurs at the cathode (Red Cat) the anode loses mass the cathode gains mass

o Explain, in terms of atoms and ions, why the anode loses mass and the cathode gains mass

o Compare and contrast voltaic cells with electrolytic cells

6.

o A voltaic cell spontaneously converts chemical energy into electrical energy.

o The purpose of the salt bridge is to allow for the migration of ions between half-cells

o Given a diagram of a voltaic cell, identify and label the cathode, anode, salt bridge, and the direction of electron flow

o Explain the function of the salt bridge and the direction of positive and negative ion migration

o Write balanced oxidation and reduction half-reactions

7.

o An electrolytic cell requires electrical energy to produce a chemical change. Electrolytic cells can be used for electrolysis (splitting a compound into its elements) and for electroplating (coating something with a metal).

o Given a diagram of an electrolytic cell, identify and label the cathode, anode, and direction of electron flow.

XI. ACIDS, BASES, AND SALTSKnowledge Application

1.

o The behavior of many acids and bases can be explained by the Arrhenius Theory.

o Arrhenius acids produce H + (hydrogen ions) as the only positive ions in aqueous solution. The hydrogen ion may also be written as H3O + and called the hydronium ion.

o Arrhenius bases produce OH – (hydroxide ions) as the only negative ion is aqueous solution.(Table E)

o Know the definitions of Arrhenius acids and bases

o If given the properties, chemical formula, or name, identify a substance as an Arrhenius acid or Arrhenius base. (Use Tables K, L, and T to help you remember these definitions. Arrhenius acids begin with H, Arrhenius bases are metals + hydroxide ion(s). *Don’t be fooled by alcohols, which also end in OH, but contain covalent bonds and do not ionize like bases do in solution. ALCOHOLS ARE NOT BASES!)

2.

o Arrhenius acids, Arrhenius bases, and salts (ionic compounds) are all electrolytes. An electrolyte is a substance which, when dissolved in water, forms a solution capable of conducting an electric current (electricity). Electrolytes can conduct electricity because they ionize (break apart into ions) in a solution.

o The ability of a solution to conduct an electric current depends on the concentration of the ions in it (more ions, more conduction).

o Given names or chemical formulas, identify acids, bases, and salts as being electrolytes

o Determine the relative strength (strong or weak) of an electrolyte given information on its ability to ionize in solution. (Strong acids and strong bases are strong electrolytes – Tables K and L list acids and bases in order from strongest to weakest. If a salt is soluble, it is a strong electrolyte – Table F can be used to determine the solubility of different salts.)

3.

o In the process of neutralization, an Arrhenius acid and an Arrhenius base react to form a salt and water. Acid + Base Salt + Water

o Recognize neutralization reactions when given the reaction equation

o Write neutralization reactions when given the reactants. (Remember that this is a double replacement reaction. Just switch the positive ions, look up their charges and cross down the subscripts if needed. Then

balance the equation.)

4.

o Titration is a laboratory process in which a volume of solution with a known concentration is added to another solution of unknown concentration. Titrations are done to determine the concentration of the unknown solution.

o Calculate the concentration or volume of a solution, using titration data using the equation Ma x Va = Mb x Vb (This equation is on Table T)

5.

o There are alternate acid-base theories. One such theory states that the acid is a proton donor (H + donor) and the base is a proton acceptor.

o Give the alternate definitions of acids and bases

o Use this definition to explain why ammonia is considered a base

6.

o The acidity or alkalinity of an aqueous solution can be measured using the pH scale.

o The pH scale measures the concentration of H+/H3O+ in a solution. [H+] = 10–pH A pH of 1 means that the [H+] = 10–1

= 0.1MA pH of 3 means that the [H+] = 10–3

= 0.001M Acids have pH values between 0 and 7 (the stronger the acid, the lower the pH and the more H+) [H+] [OH–] Neutral solutions have a pH of 7 [H+] = [OH–] Bases have pH values between 7 and 14 (the stronger the base, the higher the pH and the more OH–) [H+] [OH–]

o Identify a solution as acidic, basic (alkaline), or neutral based upon the pH value OR the relative concentrations of H+/H3O+ and OH–

o Describe acidic, basic, and neutral solutions in terms of pH value and relative H+/H3O+ and OH– concentrations

o Differentiate between strong acids/bases and weak acids/bases given pH values or ion concentrations

7.

o The pH scale is a logarithmic scale, which means that a change of one pH unit changes the concentration

o Determine the new pH value of a solution given the starting pH and the amount of increase or

of H+/H3O+ by a factor of ten tenfold = 10 times = 101 hundredfold = 100 times = 102

thousandfold = 1000 times = 103

The exponents represent the CHANGE in pH

o If a solution becomes more acidic, the pH , and the [H+]/[H3O+]

o If a solution becomes more basic, the pH , and the [H+]/[H3O+]

decrease in [H+]/[H3O+] (such as tenfold, a hundredfold, or a thousandfold)

Ex: A lake with an initial pH of 6 has been affected by acid rain. The acid rain has caused a hundredfold change in the [H+] concentration of the lake. What is the new pH of the lake?

Answer: pH = 4

o Determine the amount that the [H+]/[H3O+] would increase or decrease given a certain change in pH

8.

o The pH of a solution can be shown by using indicators.

o An indicator is a substance that changes color depending on the concentration of hydrogen/hydronium ions in a solution.

o Interpret changes in acid-base indicator color

o Explain how different indicators can be used to distinguish between solutions with different pH values

o Identify appropriate indicators that can be used to show changes in pH values, such as during a titration, given starting and ending pH values

XII. NUCLEAR CHEMISTRYKnowledge Application

1.

o The stability of an isotope is based on the ratio of the neutrons and protons in the nucleus.

o Usually when the ratio is not 1:1, the nucleus gets a little unstable and starts spitting out particles so that it will have a more stable 1:1 ratio.

o Although most nuclei are stable, some are unstable and spontaneously emit radiation. We call these unstable isotopes radioactive isotopes, radioisotopes, or nuclides.

2.

o Spontaneous decay (natural emission of radiation) by a nuclide (radioactive isotope) involves the release of particles and/or energy from the nucleus.

o Each radioactive isotope has a specific decay mode (the kind of particle or energy it gives off from its unstable nucleus) (Tables N and O!) alpha decay: release of alpha particles beta decay: release of beta particles positron emission: release of positrons gamma radiation : release of gamma rays

o These emissions differ in mass, charge, ionizing power, and penetrating power.

o Determine decay mode and write nuclear equations showing alpha decay, beta decay, positron emission, and gamma radiation (*Remember to put radioactive emissions on the RIGHT side of the arrow – if something is released, it goes on the right)

o Compare and contrast the 4 different types of radiation in terms of mass, charge, ionizing power, and penetrating power.

3.

o Each radioactive isotope has a specific half-life (rate of decay). The half-life is the time it takes for half of the radioisotope to decay/transmutate into something more stable). (Table N)

o Calculate the initial amount, the fraction remaining, time elapsed, or the half-life of a radioactive isotope, given the other variables

4.

o Nuclear reactions are represented by equations that include symbols for elements and radioactive emissions (with mass number in upper left and charge/atomic number in lower left)

o These reactions show conservation of mass and charge

o Complete nuclear equations and predict missing particles in nuclear equations

o Write nuclear equations given word problems

5.

o A change in the nucleus of an atom that changes it from one element to another is called transmutation. This can occur naturally or can be done artificially by bombarding the nucleus with high-energy particles.

o Distinguish between natural transmutation (one reactant) and artificial transmutation (two reactants) given nuclear equations

6.

o Types of nuclear reactions include fission and fusion. Fission and fusion can be natural or artificial transmutations.

o Compare and contrast fission and fusion reactions.

o Distinguish between fission and fusion reactions given nuclear equations

7 o Compare and contrast chemical

.

o Nuclear changes convert matter into energy (E = mc2)

o Energy released during nuclear reactions is much greater than the energy released during chemical reactions.

reactions and nuclear reactions o Describe benefits of using

nuclear fission

8.

o There are risks and problems associated with radioactivity and the use of radioactive isotopes, including: biological exposure, long-term storage and disposal problems, and nuclear accidents which release radioactive materials into the environment.

o Describe the risks and problems associated with using radioactive isotopes

9.

o In addition to using nuclear fission for nuclear power, radioactive isotopes have other beneficial uses in medicine and industrial chemistry, including: radioactive dating (ages of once-living things can be found from the ratio

of C-14 to C-12 in the remains; ages of rocks can be found from the ratio of U-238 to Pb-206)

tracing chemical and biological processes (radioactive tracers can be injected into the body and then x-rayed. The radioactive substance will show up on the x-ray and if there are problems, they can be detected easily)

detecting and treating of disease (Sr-90: diagnosing and treating bone cancer; I-131: diagnosing and treating thyroid disorders; Co-60: cancer treatment

radiation can be used to kill bacteria in foods (used with spices, meats, produce)