SOL Review Ms. DiOrio.

SOL Review Ms. DiOrio

Topic 1 Atomic Structure and Periodic Relationships

Matter Matter occurs as elements (pure), compounds (pure), and mixtures, which may be homogeneous (solutions) or heterogeneous. Some elements, such as oxygen, hydrogen, fluorine, chlorine, bromine, iodine, and nitrogen, naturally occur as diatomic molecules. Matter is classified by its chemical and physical properties. Physical properties refer to the condition or quality of a substance that can be observed or measured without changing the substance’s composition. Important physical properties are density, conductivity, melting point, boiling point, malleability, and ductility. Chemical properties refer to the ability of a substance to undergo chemical reaction and form a new substance. Reactivity is the tendency of an element to enter into a chemical reaction. BE ABLE TO: compare an element’s reactivity to the reactivity of other elements in the table. differentiate between pure substances and mixtures and between homogeneous and heterogeneous mixtures.

Models of the atom

Democritus ~ 430 BC (5th Century BC)
All matter is composed of eternal, indivisible, indestructible, and infinitely small substances Atomos = Greek word for indivisible Significance: This theory introduced the idea of the atom as the basic unit of matter.

John Dalton 18th – 19th Century Dalton's Atomic Theory:
Elements are made of extremely small particles called atoms. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. Atoms cannot be subdivided, created, or destroyed. Atoms of different elements combine in simple whole- number ratios to form chemical compounds In chemical reactions, atoms are combined, separated, or rearranged. Significance: Dalton’s atomic theory began a larger discussion on atomic structure that continued to grow to our current understanding. Solid Sphere Model

J.J. (Joseph John) Thomson
Discovered that cathode rays are composed of negatively charged particles smaller than protons and neutrons with a very high charge to mass ratio Introduced the “Plum Pudding” Model of the atom to include these particles Significance: Discovered the electron Plum Pudding Model

Ernest Rutherford 20th century (1909)
Discovered the concept of radioactive half-life Differentiated and names alpha and beta radiation Theorized that the charge of atoms is concentrated in the center of the atom Created the Rutherford Model Significance: Discovered the nucleus Nuclear Model

Niels Bohr Proposed discrete energy levels of electrons surrounding the nucleus Bohr Model Significance: Made foundational contributions to the understanding of atomic structure and quantum theory Planetary Model

Werner Heisenberg 1901 – 1976 Hesienberg Uncertainty Principle: the more precisely the position of some particle is determined, the less precisely its momentum can be known (and vice versa) Significance: One of the key pioneers of quantum mechanics Won a Nobel Prize in 1932 for “the creation of quantum mechanics”

Louis DeBroglie 1892 – 1987 Postulated the wave nature of electrons, suggesting that all matter has wave properties Wave-particle duality Significance: Made ground- breaking contributions to quantum theory that changed how we think about the physics of electrons

Subatomic Particles BE ABLE TO: Protons Neutrons Electrons Charge +
n/a - Mass 1 amu Location nucleus electron cloud Protons and neutrons are located in the nucleus of the atom and comprise most of its mass. Quarks are also located in the nucleus of the atom. Electrons are located in electron clouds or probability clouds outside the nucleus. BE ABLE TO: differentiate between the major atom components (proton, neutron and electron) in terms of location, size, and charge.

Atomic Mass and Isotopes
The mass number of an element is the sum of the number of protons and neutrons. It is different for each element’s isotopes. The average atomic mass for each element is the weighted average of that element’s naturally occurring isotopes. An isotope is an atom that has the same number of protons as another atom of the same element but has a different number of neutrons. Some isotopes are radioactive; many are not. BE ABLE TO: determine the number of neutrons in an isotope given its mass number. perform calculations to determine the weighted average atomic mass.

Lewis Dot Diagrams Lewis dot diagrams are used to represent valence electrons in an element. Structural formulas show the arrangements of atoms and bonds in a molecule and are represented by Lewis dot structures. BE ABLE TO: draw Lewis dot diagrams to represent valence electrons in elements and draw Lewis dot structures to show covalent bonding.

Bohr Models Bohr models show the positions of all electrons in specific energy levels

Scientists identify key contributions of principal scientists including: atomos, initial idea of atom - Democritus first atomic theory of matter, solid sphere model - John Dalton discovery of the electron using the cathode ray tube experiment, plum pudding model - J. J. Thomson discovery of the nucleus using the gold foil experiment, nuclear model - Ernest Rutherford discovery of charge of electron using the oil drop experiment - Robert Millikan energy levels, planetary model - Niels Bohr periodic table arranged by atomic mass - Dmitri Mendeleev periodic table arranged by atomic number - Henry Moseley quantum nature of energy - Max Planck uncertainty principle, quantum mechanical model - Werner Heisenberg wave theory, quantum mechanical model - Louis de Broglie.

Dmitri Mendeleev 19th Century Periodic Law:
Chemical properties go through cyclic patterns when arranged by atomic number Significance: Devised the periodic table that we use today “Father of the Periodic Table”

Henry Moseley 1887-1915 Arranged the periodic table by atomic number
Developed Moseley’s Law in X-ray spectra Moseley’s Law: sorts the chemical elements of the periodic table in a logical order based on their physics Significance: Provided evidence proving the Bohr Model of the electron

Robert Millikan 1868-1953 Oil-Drop Experiment
Tested Einstein’s Photoelectric Effect Calculated Planck’s Constant Significance: Determined the charge of an electron 1.602 x coulombs

Max Planck Early 20th Century
Rejected the classical physics model of energy as wave-like and instead assumed energy is discrete and particle-like E=hn Planck’s constant (h) = 6.62 x J/s Significance: Originator of quantum theory that changed how we think about energy and subatomic processes

Periodic Table The periodic table is arranged in order of increasing atomic numbers. The atomic number of an element is the same as the number of protons. In a neutral atom, the number of electrons is the same as the number of protons. All atoms of an element have the same number of protons. The Periodic Law states that when elements are arranged in order of increasing atomic numbers, their physical and chemical properties show a periodic pattern Periodicity is regularly repeating patterns or trends in the chemical and physical properties of the elements arranged in the periodic table. BE ABLE TO: determine the atomic number, atomic mass, the number of protons, and the number of electrons of any atom of a particular element using a periodic table.

Periodic Table The names of groups and periods on the periodic chart are alkali metals, alkaline earth metals, transition metals, halogens, and noble gases. Periods and groups are named by numbering columns and rows. Horizontal rows called periods have predictable properties based on an increasing number of electrons in the outer energy levels. Vertical columns called groups or families have similar properties because of their similar valence electron configurations. BE ABLE TO: distinguish between a group and a period. identify key groups, periods, and regions of elements on the periodic table. distinguish between physical and chemical properties of metals and nonmetals.

Periodic Trends Atomic radius is the measure of the distance between radii of two identical atoms of an element. Atomic radius decreases from left to right and increases from top to bottom within given groups. Electronegativity is the measure of the attraction of an atom for electrons in a bond. Electronegativity increases from left to right within a period and decreases from top to bottom within a group. Shielding effect is constant within a given period and increases within given groups from top to bottom. Ionization energy is the energy required to remove the most loosely held electron from a neutral atom. Ionization energies generally increase from left to right and decrease from top to bottom of a given group. BE ABLE TO: identify and explain trends in the periodic table as they relate to ionization energy, electronegativity, shielding effect, and relative sizes.

Radioactivity Half-life is the length of time required for half of a given sample of a radioactive isotope to decay. BE ABLE TO: perform calculations involving the half-life of a radioactive substance. differentiate between alpha, beta, and gamma radiation with respect to penetrating power, shielding, and composition.

Ions Atoms can gain, lose, or share electrons within the outer energy level. Loss of electrons from neutral atoms results in the formation of an ion with a positive charge (cation). Gain of electrons by a neutral atom results in the formation of an ion with a negative charge (anion).

Electron Configuration
Electron configuration is the arrangement of electrons around the nucleus of an atom based on their energy level. Electrons are added one at a time to the lowest energy levels first (Aufbau Principle). Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results (Hund’s Rule). Energy levels are designated 1-7. Orbitals are designated s, p, d, and f according to their shapes and relate to the regions of the Periodic Table. An orbital can hold a maximum of two electrons (Pauli Exclusion Principle). BE ABLE TO: relate the position of an element on the periodic table to its electron configuration. determine the number of valence electrons and possible oxidation numbers from an element’s electron configuration. write the electron configuration for the first 20 elements of the periodic table.

Orbital Diagrams 1s22s22p63s23p6
Orbital diagrams show the spin on each electron in each energy subshell 1s22s22p63s23p6

Transition Metals / Metalloids
Transition metals can have multiple oxidation states. Metalloids have properties of metals and nonmetals. They are located between metals and nonmetals on the periodic table. Some are used in semiconductors.

Redox Reactions Be able to recognize redox reactions
Redox is short for oxidation-reduction in which electrons are transferred from one species to another You can identify a redox reaction by a change in oxidation number Rules for Assigning Oxidation States The oxidation state of an element corresponds to the number of electrons that an atom loses or gains when joining with other atoms in a compound. In determine oxidation state, use the following rules: The oxidation state of an uncombined element is always 0 The total oxidation state of a neutral compound MUST equal 0 but ions equal the overall charge of the atom or molecule The oxidation state of hydrogen is generally +1 and oxygen is -2 Group 1 metals are +1 and Group 2 metals are +2 Halogens are -1, Group 6 are -2, and Group 15 are -3

Redox Reactions Fe + O2  Fe2O3 Zn + 2H+  Zn2+ + H2
Rules for Assigning Oxidation States The oxidation state of an uncombined element is always 0 The total oxidation state of a neutral compound MUST equal 0 but ions equal the overall charge of the atom or molecule The oxidation state of hydrogen is generally +1 and oxygen is -2 Group 1 metals are +1 and Group 2 metals are +2 Halogens are -1, Group 6 are -2, and Group 15 are -3 Fe (s) +O 2(g) →Fe 2 O 3(g)  F e ( s )   + O 2 ( g )   → F e 2  O 3 ( g )   Fe 2+  F e 2 +   Ag (s) +H 2 S→Ag 2 S (g) +H 2(g)  A g ( s )   + H 2  S → A g 2  S ( g )   + H 2 ( g )     Fe (s) +O 2(g) →Fe 2 O 3(g)  F e ( s )   + O 2 ( g )   → F e 2  O 3 ( g )   Fe 2+  F e 2 +   Ag (s) +H 2 S→Ag 2 S (g) +H 2(g)  A g ( s )   + H 2  S → A g 2  S ( g )   + H 2 ( g )     +3 -2 +1 +2 Fe + O2  Fe2O3 Zn + 2H+  Zn2+ + H2 SO32- + MnO4-  SO42- + Mn2+ Cu + 2Ag +  Cu2+ + 2Ag CO2 + H2  CO + H2O MnO4- + I-  I2 + Mn2+ +1 +2 +4 -2 +2 -2 +1 -2 +4 -2 +7 -2 +6 -2 +2 +7 -2 -1 +2

Zn + 2H+  Zn2+ + H2 Oxidation Reduction oxidation 0 +2 reduction +1 0
Reduction is a decrease in oxidation number Gaining electrons Oxidation is an increase in oxidation number Losing electrons Oxidation and reduction always go together. For something to be reduced, something else must be oxidized. (and vice versa) oxidation Zn + 2H+  Zn2+ + H2 reduction

LEO the lion goes GER Redox Reactions LEO = Losing electrons oxidation
GER = Gaining electrons reduction

Topic 1 Quiz Only use Part 1 for this quiz
Write down whether you have A or B for this quiz in the test record section on the front Do NOT write on quizzes (class set)

Topic 2 Chemical Formulas and Reactions

Organic Chemistry

Organic Compounds The bonding characteristics of carbon contribute to its stability and allow it to be the foundation of organic molecules. These characteristics result in the formation of a large variety of structures such as DNA, RNA, and amino acids. Carbon-based compounds include simple hydrocarbons, small carbon- containing molecules with functional groups, complex polymers, and biological molecules. Petrochemicals contain hydrocarbons Ex: propane, butane, and octane

Types of Organic Compounds
All organic compounds MUST contain carbon These compounds are given the root Alk- to identify as an organic compound The types of bonds (single, double, or triple) linking the carbons together determine the type of organic compound Alkanes = All single bonds Alkenes = Any double bonds Alkynes = Any triple bonds

Alkanes CnH2n+2 Alkanes only contain single bonds between carbon atoms
These molecules are said to be saturated because each carbon atom is surrounded by the maximum number of hydrogen atoms CnH2n+2

Alkenes Alkenes are defined as containing any double bonds between two carbon atoms (even if the rest of the bonds are single) These molecules unsaturated because the carbon atoms involved in the double bond do NOT have the maximum number of hydrogen atoms CnH2n

Alkynes Alkynes are defined as containing any triple bonds between two carbon atoms (even if the rest of the bonds are single) These molecules unsaturated because the carbon atoms involved in the double bond do NOT have the maximum number of hydrogen atoms CnH2n-2

Organic Suffix Ending C-C Bond Formula -ane Single bond CnH2n+2 -ene
Double bond CnH2n -yne Triple bond CnH2n-2

Naming Organic Compounds
Organic naming is similar to covalent naming The root of the name (originally Alk-) is changed to a prefix depending on the number of carbon atoms The suffix is determined by the type of bonds (-ane, -ene, or –yne) Number of Carbon Atoms Prefix 1 Meth- 2 Eth- 3 Prop- 4 But- 5 Pent- Number of Carbon Atoms Prefix 6 Hex- 7 Hept- 8 Oct- 9 Non- 10 Dec-

Organic Naming Practice
Number of Carbons = Types of bonds = 3 (Prop-) Propane Single (-ane)

Number of Carbons = Types of bonds = 4 (But-) Butene Double (-ene)

Number of Carbons = Types of bonds = 9 (Non-) Nonane Single (-ane)

Number of Carbons = Types of bonds = 2 (Eth-) Ethyne Triple (-yne)

Number of Carbons = Types of bonds = 5 (Pent-) Pentene Double (-ene)

Number of Carbons = Types of bonds = 1 (Meth-) Methane Single (-ane)

You do not need to memorize these!
Functional Groups Functional groups are special groups on an organic molecule that perform different functions. These functional groups get special endings to their names. Functional Group Name Example Aldehyde ______al Ethanal Alcohol _____ol Ethanol Carboxylic Acid ____oic acid Ethanoic acid Ketone _____one Propanone

Uses for Organic Compounds
Pharmaceuticals Aspirin, vitamins, and insulin are all organic molecules Plastics Plastics are formed from petrochemicals of very long carbon chains Biological Macromolecules Nucleic acids Proteins Aspirin Protein

Polymers Small molecules link to make large molecules called polymers that have combinations with repetitive subunits Natural polymers include proteins and nucleic acids Synthetic (human-made) polymers include polythene, nylon, and Kevlar

Compounds When pairs of elements form two or more compounds, the masses of one element that combine with a fixed mass of the other element form simple, whole-number ratios (Law of Multiple Proportions). Compounds have different properties than the elements from which they are composed.

Chemical Formulas NH3 H2O CO CO2 SO2 CF4
The empirical formula shows the simplest whole-number ratio in which the atoms of the elements are present in the compound. The molecular formula shows the actual number of atoms of each element in one molecule of the substance. BE ABLE TO: write the chemical formulas for certain common substances, such as ammonia, water, carbon monoxide, carbon dioxide, sulfur dioxide, and carbon tetrafluoride. NH3 H2O CO CO2 SO2 CF4 Ammonia Water Carbon monoxide Carbon dioxide Sulfur dioxide Carbon tetrafluoride

Bonding Bonds form between atoms to achieve stability.
Covalent bonds involve the sharing of electrons between atoms. Ionic bonds involve the transfer of electrons between ions. Elements with low ionization energy form positive ions (cations) easily. Elements with high ionization energy form negative ions (anions) easily.

Naming CO32- SO42- NO3- OH- PO43- NH4+
Chemical formulas are used to represent compounds. Subscripts represent the relative number of each type of atom in a molecule or formula unit. The International Union of Pure and Applied Chemistry (IUPAC) system is used for naming compounds. BE ABLE TO: name binary covalent/molecular compounds. name binary ionic compounds (using the Roman numeral system where appropriate). use polyatomic ions for naming and writing formulas of ionic compounds, including carbonate, sulfate, nitrate, hydroxide, phosphate, and ammonium. CO32- SO42- NO3- OH- PO43- NH4+ carbonate sulfate nitrate hydroxide phosphate ammonium

VSEPR Theory predict, draw, and name molecular shapes (bent, linear, trigonal planar, tetrahedral, and trigonal pyramidal). BE ABLE TO: use valence shell electron pair repulsion (VSEPR) model to draw and name molecular shapes (bent, linear, trigonal planar, tetrahedral, and trigonal pyramidal).

Polarity Polar molecules result when electrons are distributed unequally. Polar bonds form between elements with very different electronegativities. Non-polar bonds form between elements with similar electronegativities. BE ABLE TO: recognize polar molecules and non-polar molecules. : : : : : : : : : :

Types of Chemical Reactions
Major types of chemical reactions are synthesis (A+B  AB) decomposition (BC  B+C) single replacement (A+BC  B+AC) double replacement (AC+BD  AD+BC) neutralization (HX+MOH  H2O + MX) combustion (CxHy + O2  CO2 + H2O). BE ABLE TO: classify types of chemical reactions as synthesis, decomposition, single replacement, double replacement, neutralization, and/or combustion. recognize equations for redox reactions and neutralization reactions.

Balancing 2H2 + O2  2H2O H-H O-O H-H-O H-H H-H-O
Conservation of matter is represented in balanced chemical equations. A coefficient is a quantity that precedes a reactant or product formula in a chemical equation and indicates the relative number of particles involved in the reaction. 2H2 + O2  2H2O H-H O-O H-H-O H-H H-H-O Element Reactants Products H O

Writing Equations BE ABLE TO:
transform word equations into chemical equations and balance chemical equations.

Kinetics Kinetics is the study of reaction rates. Several factors affect reaction rates, including temperature, concentration, surface area, and the presence of a catalyst. Reaction rates/kinetics are affected by activation energy, catalysis, and the degree of randomness (entropy). Catalysts decrease the amount of activation energy needed. BE ABLE TO: interpret reaction rate diagrams. identify and explain the effect the following factors have on the rate of a chemical reaction: catalyst, temperature, concentration, size of particles.

Exothermic vs. Endothermic
Chemical reactions are exothermic reactions (heat producing) and endothermic reactions (heat absorbing). BE ABLE TO: distinguish between an endothermic and exothermic process.

Equilibrium Reactions occurring in both forward and reverse directions are reversible. Reversible reactions can reach a state of equilibrium, where the reaction rates of both the forward and reverse reactions are constant. Le Chatelier’s Principle indicates the qualitative prediction of direction of change with temperature, pressure, and concentration. BE ABLE TO: recognize that there is a natural tendency for systems to move in a direction of randomness (entropy). predict the shift in equilibrium when a system is subjected to a stress (Le Chatelier’s Principle) and identify the factors that can cause a shift in equilibrium (temperature, pressure, and concentration.) distinguish between irreversible reactions and those at equilibrium.

Topic 2 Warm-Up Identify the following types of reactions:
A + B  AB A + BC  B + AC CH4 + 2O2  CO2 + 2H2O Write the formula for the following compounds: Copper (II) oxide Sodium sulfide Name the following compounds: MgO CCl4 Na2SO4 List the three most common synthetic organic polymers.

Topic 1 Most Missed Questions

Topic 1: Most Missed Questions
2. Compared to atoms of metals, atoms of nonmetals generally have Higher electronegativities and higher ionization energies Lower electronegativities and higher ionization energies Lower electronegativities and lower ionization energies Higher electronegativities and lower ionization energies Metals Nonmetals

6. Alkali metals, alkaline earth metals, and halogens are found respectively in Groups 2, 13, and 17 1, 2, and 17 1, 2, and 18 1, 2, and 14

8. What is the oxidation number of chlorine in HClO4? +3 +1 +7 +5 HClO4

8. What is the oxidation number of chlorine in HClO4? +3 +1 +7 +5 -2 +1 HClO4 1 + x – 8 = 0

10. In the reaction Mg + 2HCl  MgCl2 + H2, the magnesium Gains electrons and is reduced Loses electrons and is reduced Loses electrons and is oxidized Gains electrons is oxidized +2 Mg + 2HCl  MgCl2 + H2 LEO the lion goes GER

15. The atoms of which element require the greatest amount of energy to remove an electron? Neon Helium Krypton Argon

19. Which element forms a diatomic molecule containing a triple covalent bond? Chlorine Hydrogen Oxygen Nitrogen

24. In the diagram below, the radiation from a radioactive source is being separated as it passes between electrically charged plates. What are the three types of radiation observed on the detector? X = gamma, Y = beta, Z = alpha X = beta, Y = gamma, Z = alpha X = alpha, Y = beta, Z = gamma X = gamma, Y = alpha, Z = beta

Write down whether you have A or B for this quiz in the test record section on the front Do NOT write on quizzes (class set) NAME: Ms. DiOrio #0 SUBJECT: Chem Topic 2-5 Quizzes Test Record Part 1 A or B Part 2

Topic 3 Phases of Matter and Kinetic Molecular theory

Phase of Matter The phase of a substance depends on temperature and pressure. Forces of attraction (intermolecular forces) between molecules determine their state of matter at a given temperature. Solid, liquid, and gas phases of a substance have different energy content. Pressure, temperature, and volume changes can cause a change in physical state. Specific amounts of energy are absorbed or released during phase changes. A fourth phase of matter is plasma. Plasma is formed when a gas is heated to a temperature at which its electrons dissociate from the nuclei.

Intermolecular Forces
Forces of attraction include hydrogen bonding, dipole-dipole attraction, and London dispersion (van der Waals) forces. BE ABLE TO: identify how hydrogen bonding in water plays an important role in many physical, chemical, and biological phenomena.

Phase Changes Solid  liquid = Liquid  gas = Solid  gas =
Gas  liquid = Liquid  solid = Gas  solid = Melting Vaporization Sublimation Condensation Freezing Deposition

Heating Curve A heating curve graphically describes the relationship between temperature and energy (heat). It can be used to identify a substance’s phase of matter at a given temperature as well as the temperature(s) at which it changes phase. It also shows the strength of the intermolecular forces present in a substance. BE ABLE TO: graph and interpret a heating curve (temperature vs. time).

Q = mCΔT Q = ΔHphase change(moles or grams)
Calorimetry Molar heat of fusion is a property that describes the amount of energy needed to convert one mole of a substance between its solid and liquid states. Molar heat of vaporization is a property that describes the amount of energy needed to convert one mole of a substance between its liquid and gas states. Specific heat capacity is a property of a substance that tells the amount of energy needed to raise one gram of a substance by one degree Celsius. The values of these properties are related to the strength of their intermolecular forces. Q = mCΔT Q = ΔHphase change(moles or grams) BE ABLE TO: calculate energy changes, using molar heat of fusion and molar heat of vaporization. calculate energy changes, using specific heat capacity.

Phase Diagrams Triple point – All three phases exist in equilibrium
Critical point – A liquid can no longer exist beyond this temperature/pressure BE ABLE TO: interpret a phase diagram of water.

Properties of Liquids Vapor pressure is the pressure of the vapor found directly above a liquid in a closed container. When the vapor pressure equals the atmospheric pressure, a liquid boils. Volatile liquids have high vapor pressures, weak intermolecular forces, and low boiling points. Nonvolatile liquids have low vapor pressures, strong intermolecular forces, and high boiling points. BE ABLE TO: interpret vapor pressure graphs.

Solutions Solutions can be a variety of solute/solvent combinations: gas/gas, gas/liquid, liquid/liquid, solid/liquid, gas/solid, liquid/solid, or solid/solid.

Solubility The saturation of a solution is dependent on the amount of solute present in the solution. Polar substances dissolve ionic or polar substances; nonpolar substances dissolve nonpolar substances. BE ABLE TO: interpret solubility curves.

Colligative Properties
A liquid’s boiling point and freezing point are affected by changes in atmospheric pressure. A liquid’s boiling point and freezing point are affected by the presence of certain solutes. The number of solute particles changes the freezing point and boiling point of a pure substance.

Kinetic Molecular Theory
The kinetic molecular theory is a model for predicting and explaining gas behavior. Atoms and molecules are in constant motion. Temperature is a measurement of the average kinetic energy in a sample. There is a direct relationship between temperature and average kinetic energy.

Ptotal = P1 + P2 + … Ppartial = X(Ptotal)
Gases Gases have mass and occupy space. Gas particles are in constant, rapid, random motion and exert pressure as they collide with the walls of their containers. Gas molecules with the lightest mass travel fastest. Relatively large distances separate gas particles from each other. Equal volumes of gases at the same temperature and pressure contain an equal number of particles. (Avogadro’s Principle) Pressure units include atm, kPa, and mm Hg. An ideal gas does not exist, but this concept is used to model gas behavior. A real gas exists, has intermolecular forces and particle volume, and can change states. The Ideal Gas Law states that PV = nRT. The sum of the partial pressures of all the components in a gas mixture is equal to the total pressure of a gas mixture (Dalton’s law of partial pressures). Ptotal = P1 + P2 + … Ppartial = X(Ptotal)

Gas Laws The pressure and volume of a sample of a gas at constant temperature are inversely proportional to each other (Boyle’s Law: P1V1 = P2V2). At constant pressure, the volume of a fixed amount of gas is directly proportional to its absolute temperature (Charles’ Law: V1/T1 = V2/T2). The Combined Gas Law (P1V1/T1 = P2V2/T2) relates pressure, volume, and temperature of a gas. BE ABLE TO: explain the behavior of gases and the relationship between pressure and volume (Boyle’s Law), and volume and temperature (Charles’ Law). solve problems and interpret graphs involving the gas laws.

Standard Temperature Pressure
0oC 1 atm 760 torr 273 K 760 mmHg 101.3 kPa

34. What is the empirical formula of a compound that contains 92.3% carbon and 7.7% hydrogen by mass? CH4 CH2 CH3 CH 92.3 g C 1 mol 12.01 g = 7.69 mol C / 7.62 = 1 mol C 7.7 g H 1 mol 1.01 g = 7.62 mol H / 7.62 = 1 mol H

34. What is the empirical formula of a compound that contains 92.3% carbon and 7.7% hydrogen by mass? CH4 CH2 CH3 CH 92.3 g C 1 mol 12.01 g = 7.69 mol C / 7.62 = 1 mol C CH 7.7 g C 1 mol 1.01 g = 7.62 mol H / 7.62 = 1 mol H

39. Given the reaction at equilibrium: A2(g) + B2(g) ⇋ 2AB(g) + heat Which stress on the system at equilibrium will increase the concentration of AB(g)? Increasing the temperature Decreasing the concentration of A2(g) Decreasing the pressure Increasing the concentration of B2(g)

40. Given the reaction at equilibrium: N2(g) + 3H2(g) ⇋ 2NH3(g) + heat Which change would shift the equilibrium to the right? Decrease the [H2] Increase the temperature Increase the pressure Decrease the [N2]

42. Which electron-dot formula represents a polar molecule? H : H g F f H : C : H H : : : : F : : : H : Cl :

48. What is the formula for ammonium carbonate? (NH4)2(CO3)2 (NH4)2CO3 NH4(CO3)2 NH4CO3 NH4+ CO32-

48. What is the formula for ammonium carbonate? (NH4)2(CO3)2 (NH4)2CO3 NH4(CO3)2 NH4CO3 NH4+ CO32- (NH4)2CO3

Topic 3 Warm-Up What is the difference between a real and an ideal gas? Average kinetic energy is __________ proportional to temperature. What are the three colligative properties? How are pressure and volume related? Whose law states this relationship? When do you use Q=mCΔT vs. Q= ΔHpc(mass or moles) ?

Write down whether you have A or B for this quiz in the test record section on the front Do NOT write on quizzes (class set) NAME: Ms. DiOrio #0 SUBJECT: Chem Topic 2-5 Quizzes Test Record Part 1 Part 2 A or B

Topic 4 Molar Relationships

The Mole Atoms and molecules are too small to count by usual means. A mole is a way of counting any type of particle (atoms, molecules, and formula units). Avogadro’s number = 6.02 × 10²³ particles per mole. BE ABLE TO: calculate mole ratios, percent composition, conversions, and average atomic mass.

Molar Conversions Molar mass of a substance is its average atomic mass in grams from the Periodic Table Molar volume = 22.4 L/mole for any gas at standard temperature and pressure (STP). BE ABLE TO: perform conversions between mass, volume, particles, and moles of a substance.

Stoichiometry Stoichiometry involves quantitative relationships. Stoichiometric relationships are based on mole quantities in a balanced equation. Total grams of reactant(s) = total grams of product(s). BE ABLE TO: perform stoichiometric calculations involving the following relationships: mole-mole; mass-mass; mole-mass; mass-volume; mole-volume; volume-volume; mole-particle; mass-particle; and volume-particle.

Volume A Volume B Mass A Moles A Moles B Mass B Particles A
Molar Volume 1 mol = 22.4 L Molar Volume 1 mol = 22.4 L Mass A Moles A Moles B Mass B Molar Mass A 1 mol = ___ g Molar Ratio __ mol A= __ mol B Molar Mass B 1 mol = ___ g Avogadro’s Number 1 mol = 6.02 x 1023 partices Avogadro’s Number 1 mol = 6.02 x 1023 partices Particles A Particles B

Limiting Reactant and % Yield
BE ABLE TO: identify the limiting reactant (reagent) in a reaction. calculate percent yield of a reaction.

Concentration Molarity = moles of solute/L of solution.
[ ] refers to molar concentration. When solutions are diluted, the moles of solute present initially remain. M1V1 = M2V2 BE ABLE TO: perform calculations involving the molarity of a solution, including dilutions.

Electrolytes Strong electrolytes dissociate completely. Weak electrolytes dissociate partially. Non- electrolytes do not dissociate. Non-electrolytes cannot conduct electricity Salts and strong acids/bases are strong electrolytes Weak acids and bases are weak electrolytes Covalent compounds (like sugar C6H12O6) are non-electrolytes BE ABLE TO: compare and contrast the differences between strong, weak, and non-electrolytes.

Acids and Bases Two important classes of compounds are acids and bases. Acids and bases are defined by several theories: According to the Arrhenius theory, acids are characterized by their sour taste, low pH, and the fact that they turn litmus paper red. According to the Arrhenius theory, bases are characterized by their bitter taste, slippery feel, high pH, and the fact that they turn litmus paper blue. According to the Bronsted-Lowry theory, acids are proton donors, whereas bases are proton acceptors. Acids and bases dissociate in varying degrees. BE ABLE TO: identify common examples of acids and bases, including vinegar and ammonia. differentiate between the defining characteristics of the Arrhenius theory of acids and bases and the Bronsted-Lowry theory of acids and bases.

pH pH is a number scale ranging from 0 to 14 that represents the acidity of a solution. The pH number denotes hydrogen (hydronium) ion concentration. The pOH number denotes hydroxide ion concentration. The higher the hydronium [H3O+] concentration, the lower the pH. pH + pOH = 14 BE ABLE TO: relate the hydronium ion concentration to the pH scale.

Titration Strong acid-strong base titration is the process that measures [H+] and [OH - ]. Indicators show color changes at certain pH levels. BE ABLE TO: perform titrations in a laboratory setting using indicators.

57. Which of these values is most responsible for changing the boiling and freezing points of a solvent? Number of solute particles Electronegativity of the solvent Molar mass of the solvent Weight of the solute particles Colligative properties depend only on the number of solute particles

58. Which substance is made of particles with the highest average kinetic energy? Fe(s) at 35˚C Br2(l) at 20˚C H2O(l) at 30˚C CO2(g) at 25˚C Temperature is directly proportional to average kinetic energy.

59. What interval represents the heat of reaction (ΔH)? A B C D

60. Interval B represents the Activation energy Activated complex Potential energy of the produces Potential energy of the reactants

75. The graph below represents the relationship between temperature and time for a substance that was heated uniformly starting at t0. The substance was in the solid phase at t0. During what time interval does the heat absorbed by the substance represent the heat of fusion of the substance? t1 to t2 t0 to t1 t2 to t3 t3 to t4 t1 t2 t3 t4

Topic 4 Warm - Up How many molecules are present in 2 moles of nitrogen gas? What is the gram formula mass (molar mass) of CO2? What is the percent by mass of oxygen in CO2? What is the concentration of 2.0 moles of NaCl in 500 mL of water? What are the five strong acids? 2 mol 6.02 x 1023 molecules 1 mol = 1.2 x 1024 molecules (16.00) = g/mol 2(16.00) 44.01 x 100% = 72.7% 2.0 mol 0.5 L = 4 M HCl, HBr, HI, HNO3, H2SO4

Write down whether you have A or B for this quiz in the test record section on the front Do NOT write on quizzes (class set) NAME: Ms. DiOrio #0 SUBJECT: Chem Topic 2-5 Quizzes Test Record Part 3 A or B Part 4

Topic 5 Scientific Investigation

The Nature of Science The nature of science refers to the foundational concepts that govern the way scientists formulate explanations about the natural world. The nature of science includes the following concepts the natural world is understandable; science is based on evidence - both observational and experimental; science is a blend of logic and innovation; scientific ideas are durable yet subject to change as new data are collected; science is a complex social endeavor; and scientists try to remain objective and engage in peer review to help avoid bias.

The Nature of Science Constant reevaluation in the light of new data is essential to keeping scientific knowledge current. In this fashion, all forms of scientific knowledge remain flexible and may be revised as new data and new ways of looking at existing data become available. Scientific questions drive new technologies that allow discovery of additional data and generate better questions. New tools and instruments provide an increased understanding of matter at the atomic, nano, and molecular scale. make connections between components of the nature of science and their investigations and the greater body of scientific knowledge and research.

Measurement Measurements are useful in gathering data about chemicals and how they behave. Data tables are used to record and organize measurements. Measurements of quantity include length, volume, mass, temperature, time, and pressure to the correct number of significant digits. Measurements must be expressed in International System of Units (SI) units. The last digit of any valid measurement must be estimated and is therefore uncertain. use common SI prefixes and their values (milli-, centi-, kilo-) in measurements and calculations. read a measurement from a graduated scale, stating measured digits plus the estimated digit. read measurements and record data, reporting the significant digits of the measuring equipment. make the following measurements, using the specified equipment: volume: graduated cylinder, volumetric flask, buret mass: triple beam and electronic balances temperature: thermometer and/or temperature probe pressure: barometer and/or pressure probe.

Experimental Design Repeated trials during experimentation ensure verifiable data. design and perform controlled experiments to test predictions, including the following key components: hypotheses, independent and dependent variables, constants, controls, and repeated trials. predict outcome(s) when a variable is changed. discover and eliminate procedural errors. summarize knowledge gained through gathering and appropriate processing of data in a report that documents background, objective(s), data collection, data analysis and conclusions.

Data Analysis Mathematical procedures are used to validate data, including percent error to evaluate accuracy. Algebraic equations represent relationships between dependent and independent variables. Graphing calculators can be used to manage the mathematics of chemistry. Ratios and proportions are used in calculations. determine the mean of a set of measurements. use graphing calculators to solve chemistry problems. use data collected to calculate percent error. use appropriate technology for data collection and analysis, including probeware interfaced to a graphing calculator and/or computer and computer simulations.

Sig Figs and Scientific Notation
Significant digits of a measurement are the number of known digits together with one estimated digit. Scientific notation is used to write very small and very large numbers. demonstrate the use of scientific notation, using the correct number of significant digits with powers of ten notation for the decimal place. perform calculations according to significant digits rules.

Dimensional Analysis Dimensional analysis is a way of translating a measurement from one unit to another unit. convert measurements using dimensional analysis.

Lab Techniques Techniques for experimentation involve the identification and the proper use of chemicals, the description of equipment, and the recommended statewide framework for high school laboratory safety. demonstrate the following basic lab techniques: filtering, using chromatography, and lighting a gas burner. demonstrate safe laboratory practices, procedures, and techniques.

Safety understand Material Safety Data Sheet (MSDS) warnings, including handling chemicals, lethal dose (LD), hazards, disposal, and chemical spill cleanup. identify, locate, and know how to use laboratory safety equipment, including aprons, goggles, gloves, fire extinguishers, fire blanket, safety shower, eye wash, broken glass container, and fume hood.

Lab Equipment identify the following basic lab equipment: beaker, Erlenmeyer flask, graduated cylinder, test tube, test tube rack, test tube holder, ring stand, wire gauze, clay triangle, crucible with lid, evaporating dish, watch glass, wash bottle, and dropping pipette.

Precision vs. Accuracy demonstrate precision (reproducibility) in measurement. recognize accuracy in terms of closeness to the true value of a measurement.

Graphing Graphs are used to summarize the relationship between the independent and dependent variable. Graphed data give a picture of a relationship. graph data utilizing the following: independent variable (horizontal axis) dependent variable (vertical axis) scale and units of a graph regression line (best fit curve). DV IV

Atomic Theory explain the emergence of modern theories based on historical development. For example, students should be able to explain the origin of the atomic theory beginning with the Greek atomists and continuing through the most modern quantum models. differentiate between the historical and quantum models of the atom.

108. Given the reaction: 2Na + 2H2O  2NaOH + H2 How many moles of Na are needed to produce 5.6 liters of H2 measured at STP? 1.0 2.0 0.25 0.50 1 𝑚𝑜𝑙 2 𝑚𝑜𝑙 𝑁𝑎 5.6 𝐿 𝐻2 𝑥 𝑥 = 22.4 𝐿 1 𝑚𝑜𝑙 𝐻2

112. What is the molarity of a solution of KNO3 (molecular mass – 101 grams) that contains 404 grams of KNO3 in 2.00 liters of solution? 0.50 M 1.00 M 2.00 M 4.00 M 𝐶𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛= 𝑚𝑜𝑙𝑒𝑠 𝐿𝑖𝑡𝑒𝑟𝑠 = = 4.00 𝑚𝑜𝑙 2.00 𝐿

113. Which compound is a strong electrolyte? CH3COOH HF C6H12O6 HNO3

119. How many grams of sodium chloride are required to prepare mL of a M solution? 1.46 g 2.93 g 29.3 g 58.5 g 0.100 𝑀= 𝑥 𝐿 𝑥=0.05 𝑚𝑜𝑙 𝑁𝑎𝐶𝑙 58.44 𝑔 0.05 𝑚𝑜𝑙 𝑁𝑎𝐶𝑙 𝑥 = 1 𝑚𝑜𝑙

125. Which of the following is the solute in CaCl2(aq)? Ca2+ CaCl2 Cl- H2O

Topic 5 Warm-Up 3.7 60. cm3 D = mass / volume X axis 3.5 mL
What is the sum of to the correct number of sig figs? Calculate the volume of a cube that is 4.0 cm x 3.0 cm x 5.0 cm. What is the formula for density? On a graph, the independent variable goes on which axis? What is the volume in the graduated cylinder to the right? 3.7 60. cm3 D = mass / volume X axis 3.5 mL

Write down whether you have A or B for this quiz in the test record section on the front Do NOT write on quizzes (class set) NAME: Ms. DiOrio #0 SUBJECT: Chem Topic 2-5 Quizzes Test Record Part 3 Part 4 A or B

2. f f

SOL Review Ms. DiOrio.

Similar presentations

Presentation on theme: "SOL Review Ms. DiOrio."— Presentation transcript:

Similar presentations

About project

Feedback

Log in

Auth with social network:

SOL Review Ms. DiOrio.

Similar presentations

Presentation on theme: "SOL Review Ms. DiOrio."— Presentation transcript:

Similar presentations

About project

Feedback