Is Chemistry Required in High School in Pennsylvania?
High school students in Pennsylvania are given a range of options and pathways to graduate with the individualized skills they need to pursue their chosen careers. Instead of credits, Pennsylvania opts for a performance-based system. The pathways that demonstrate a student’s readiness to graduate include, but are not limited to:
- Students achieving a minimum scaled score of 1500 or better in each of the three Keystone Exams (Algebra I, Biology, and Literature) demonstrate Keystone Proficiency and meet statewide requirements for high school graduation.
- Students who do not have a score in all three Keystone Exam areas do not qualify for either the Keystone Proficiency or Composite Pathways. Instead, they must meet statewide graduation requirements under:
- The CTE Concentrator Pathway;
- Alternative Assessment Pathway; OR
- Evidence-Based Pathway
The typical high school student in Pennsylvania will explore the following chemistry topics in preparation for their particular graduation pathway:
Chemistry: Term 1 Unit 1
Topic: Matter and Energy
- CHEM.A.1.1.1: Classify physical or chemical changes within a system in terms of matter and/or energy.
- CHEM.A.1.1.2: Classify observations as qualitative or quantitative.
- CHEM.A.1.1.3: Utilize significant figures to communicate the uncertainty in a quantitative observation.
- CHEM.A.1.1.4: Relate the physical properties of matter to is atomic or molecular structure.
- CHEM.A.1.2.1: Compare properties of solutions containing ionic or molecular solutes (e.g. dissolving, dissociating).
- CHEM.A.1.2.2: Differentiate between homogenous and heterogeneous mixtures (e.g. how such mixtures can be separated).
- CHEM.A.1.2.3: Describe how factors (e.g. temperature, concentration, surface area) can affect solubility.
- CHEM.A.1.2.4: Describe various ways that concentration can be expressed and calculated.
- CHEM.A.1.2.5: Describe how chemical bonding can affect whether a substance dissolves in a given liquid.
Key Terms and Definitions Students Will Learn
1. Allotrope - molecules of an element formed from different numbers of the same type of atom
2. atom - smallest unit of element that maintains properties of that element
3. boiling point - the temperature and pressure at which a liquid and a gas are in equilibrium
4. chemical - any substance that has a defined composition
5. chemical change - when 1 or more substances change into entirely new substances with new properties
6. chemical property - a property of matter that describes a substance’s ability to participate in chemical reactions
7. chemical reactions - process when one or more substances change to produce one or more different substances
8. compound - substance made of atoms of 2 or more different elements joined by chemical bonds
9. conservation of matter - a fundamental principle stating that matter cannot be created or destroyed in ordinary chemical and physical changes
10. Endothermic - process where heat is input into a system from the environment
11. exothermic - process where heat is released from a system into the environment
12. element - pure substances that contain only one type of atom
13. freezing point -the temperature at which a solid and liquid are in equilibrium at 1 atm pressure; the temperature at which a liquid substance freezes
14. heterogeneous - not uniform; dissimilar components; often varying in size or naturally separating in time
15. homogeneous - uniform structure or composition throughout
16. mass - quantity of matter in an object (SI unit = kg)
17. matter - anything that has mass and volume
18. melting point - the temperature and pressure at which a solid becomes a liquid
19. mixture - combination of 2 or more substances that are not chemically bonded
20. molecule - smallest unit of a substance that keeps all of the physical and chemical properties of that substance; it can consist of 2 or more atoms bonded
21. physical property - a defining characteristic of a pure substance (density, color, hardness, melting point)
22. products - substances on the right-hand side of the arrow; formed in a reaction
23. reactants - substances on the left-hand side of the arrow; used up in a reaction
24. volume - space an object occupies
25. weight - force of gravity on mass; expressed in newtons
26. Accuracy - describes how close a measurement is to the actual value.
27. Precision - exactness or reproducibility of a measurement.
28. Significant figures - limiting the precision of a reported measurement based on equipment used or mathematical operations performed.
Chemistry: Term 1 Unit 2
Topic: Atomic Structure
- CHEM.A.2.1.1: Describe the evolution of atomic theory leading to the current model of the atom based on the works of Dalton, Thomson, Rutherford, and Bohr.
- CHEM.A.2.1.2: Differentiate between the mass number of an isotope and the average atomic mass of an element.
- CHEM.A.2.2.1: Predict the ground state electron configuration and/ or orbital diagram for a given atom or ion.
- CHEM.A.2.2.3: Explain the relationship between the electron configuration and the atomic structure of a given atom or ion (e.g. energy levels and/ or orbitals with electrons, distribution of electrons in orbitals, shapes of orbitals).
- CHEM.A.2.2.4: Relate the existence of quantized energy levels to atomic emission spectra.
- CHEM.B.1.1.1: Apply the mole concept to representative particles (e.g., counting, determining mass of atoms, ions, molecules, and/or formula units).
- CHEM.B.1.2.2: Apply the law of definite proportions to the classification of elements and compounds as pure substances.
Key Terms and Definitions Students Will Learn
1. atom - smallest unit of element that maintains properties of that element
2. atomic theory - all matter is composed of discrete units called atoms
3. atomic mass - the mass of an atom expressed in atomic mass units
4. atomic number - the number of protons in the nucleus of an atom; the atomic number is the same for all atoms of an element; this number defines the element.
5. electron configuration - the arrangement of electrons in an atom
6. electron - a subatomic particle that has a negative charge
7. excited state - a state in which an atom has more energy than it does at its ground state
8. ground state - the lowest energy state of a quantized system
9. isotope - an atom that has the same number of protons (atomic number) as other atoms of the same element do but that has a different number of neutrons
10. mass number - the sum of the numbers of protons and neutrons in the nucleus of an atom
11. neutron - a subatomic particle that has no charge and that is found in the nucleus of an atom
12. nucleus - in physical science, an atom’s central region, which is made up of protons and neutrons
13. orbital - a region in an atom where there is a high probability of finding electrons
14. proton - a subatomic particle that has a positive charge and that is found in the nucleus of an atom; the number of protons of the nucleus is the atomic number, which determines the identity of an element
15. Law of definite proportions - a chemical compound always contains the same elements in exactly the same proportions by weight or mass.
16. Law of conservation of mass - mass cannot be created nor destroyed; only transformed.
17. Law of multiple proportions - when two elements combine to form two or more compounds, the mass of one element that combines with a given mass of the other is in the ratio of small whole numbers.
18. Coulomb’s law - the closer two charges are, the greater the force between them.
19. Quantum number - one of four values that defines the properties of an electron.
20. Pauli exclusion principle - Two particles of a certain class cannot be in the exact same energy state.
21. Aufbau principle - The structure of each successive element is obtained by adding one proton to the nucleus of the atom and one electron to the lowest-energy orbital that is available.
22. Hund’s rule - For an atom in the ground state, the number of unpaired electrons is the maximum possible and these unpaired electrons have the same spin.
23. Mole - The SI base unit used to measure the amount of a substance whose number of particles is the same as the number of atoms in 12 g of carbon-12.
24. Molar Mass - The mass in grams of 1 mole of a substance.
25. Avogadro’s number - 6.022 X 1023, the number of particles in one mole
Chemistry: Term 1 Unit 3
Topic: Periodic Table and Periodic Trends
- CHEM.A.2.2.2: Predict characteristics of an atom or an ion based on its location on the periodic table (e.g. number of valence electrons, potential types of bonds, reactivity).
- CHEM.A.2.2.3: Explain the relationship between the electron configuration and the atomic structure of a given atom or ion (e.g. energy levels and/or orbitals with electrons, distribution of electrons in orbitals, shapes of orbitals).
- CHEM.A.2.3.1: Explain how the periodicity of chemical properties led to the arrangement of elements on the periodic table.
- CHEM.A.2.3.2: Compare and/or predict the properties (e.g. electron affinity, ionization energy, chemical reactivity, electronegativity, atomic radius) of selected elements by using their locations on the periodic table and known trends.
Key Terms and Definitions Students Will Learn
1. periodic law - repeating chemical and physical properties of elements change periodically with the atomic numbers of the elements
2. periodicity - synonym for periodic law
3. Law of octaves - periodic law referring to the main-group’s 8 valence electrons and repeating chemical properties.
4. valence electron - an electron in the outermost energy level ( s and p shell; highest energy level only )
5. group - a vertical column of elements in the periodic table; elements in a group share chemical properties and have the same number of valence electrons.
6. chemical property - a property of matter that describes a substance’s ability to participate in chemical reactions
7. period - in chemistry, a horizontal row of elements in the periodic table
8. Atomic radius - distance from the center of the nucleus to the edge of the electron cloud.
9. bond radius - One way of measuring the atomic radius. One-half the distance between the nuclei of identical atoms that are bonded together
10. Ionic Radius - atomic radius of an ion (cation or anion).
11. alkali metal - one of the elements of Group 1 of the periodic table (lithium, sodium, potassium, rubidium, cesium, and francium)
12. alkaline-earth metal - one of the elements of Group 2 of the periodic table (beryllium, magnesium, calcium, strontium, barium, and radium)
13. halogen - one of the elements of Group 17 (fluorine, chlorine, bromine, iodine, and astatine); halogens combine with most metals to form salts
14. noble gas - a Group 18 element (helium, neon, argon, krypton, xenon, and radon). Also known as inert gases.
15. transition metal - one of the metals located in Groups 3-12 and the lower portion of the p-block.
16. Main Group - Elements from groups 1-2 and 13-18. Representing a majority of chemical properties.
17. lanthanide series- a member of the rare-earth series of elements, whose atomic numbers range from 58 (cerium) to 71 (lutetium)
18. actinide series- any of the elements of the actinide series, which have atomic numbers from 89 (actinium, Ac) through 103 (lawrencium, Lr)
19. ionization energy - the amount of energy needed to remove a single electron from an element in its gaseous state
20. electronegativity - a measure of the ability of an atom in a chemical compound to attract electrons from within the compound.
21. Metalloids - step elements defining the boundary between metals and nonmetals. Sometimes referred to as semi-conductors.
22. Electron Affinity - ability of an atom to pull electron towards itself from a different compound.
23. Electron shielding - inner electrons blocking the charge differential between the nucleus and outer electrons.
24. Alloy - a mixture of more than one metal; a solution of metals
25. Inert - having limited or no chemical reactivity.
26. Malleable - property of a metal to be hammered into sheets.
27. Ductile - Property of a metal to be pulled into a wire.
28. Nuclear reaction - reaction in which the atomic number and / or mass changes.
29. Super heavy elements - Elements which atomic numbers larger than 92. All are unstable and undergo radioactive decay.
Chemistry: Term 2 Unit 4
Topic: Ionic Bonding
- CHEM.A.1.2.5 Describe how chemical bonding can affect whether a substance dissolves in a given liquid
- CHEM.B.1.3.3 Use illustrations to predict the polarity of a molecule.
- CHEM.A.1.1.4 Relate the physical properties of matter to its atomic or molecular structure.
- CHEM.A.1.1.5 Apply a systematic set of rules (IUPAC) for naming compounds and writing chemical formulas (e.g., binary covalent, binary ionic, ionic compounds containing polyatomic ions)
- CHEM.A.2.2.2 Predict characteristics of an atom or an ion based on its location on the periodic table (e.g., number of valence electrons, potential types of bonds, reactivity).
- CHEM.A.2.3.2 Compare and/or predict the properties (e.g., electron affinity, ionization energy, chemical reactivity, electronegativity, atomic radius) of selected elements by using their locations on the periodic table and known trends.
- CHEM.B.1.2.1 Determine the empirical and molecular formulas of compounds.
- CHEM.B.1.2.2 Apply the law of definite proportions to the classification of elements and compounds as pure substances.
- CHEM.B.1.4.1 Recognize and describe different types of models that can be used to illustrate the bonds that hold atoms together in a compound (e.g., computer models, ball-and-stick models, graphical models, solid-sphere models, structural formulas, skeletal formulas, Lewis dot structures).
Key Terms and Definitions Students Will Learn
1. anion: an ion that has a negative charge
2. cation: an ion that has a positive charge
3. crystal lattice: the regular pattern in which an atom or molecules are arranged in the solid state.
4. Unit cell: smallest repetitive unit of a crystal lattice.
5. Lattice energy: the energy associated with constructing a crystal lattice relative to the energy of all constituent atoms separated by infinite distances.
6. ion: an atom, radical, or molecule that has gained or lost one or more electrons and has a negative or positive charge
7. law of definite proportions: the law that states that a chemical compound always contains the same elements in exactly the same proportions by weight or mass
8. octet rule: a concept of chemical bonding theory that is based on the assumption that atoms tend to have either empty valence shells or full valence shells of eight electrons
9. Noble gas configuration: refers to the full octet of the noble gases and a stable electronic conformation.
10. polyatomic ion: an ion made of two or more atoms; typically covalent in nature and do not dissociate when dissolved
11. salt: an ionic compound that forms when a metal atom or a positive radical replaces the hydrogen of an acid
12. valence electron: electrons in the outermost energy level of an atom
13. bond energy: a measure of the amount of energy needed to break apart one mole of covalently bonded gases. The SI units used to describe bond energy is kiloJoules per mole of bonds (kJ/mol).
Chemistry: Term 2 Unit 5
Topic: Covalent Bonds and Nomenclature
- CHEM.A.1.2.5 Describe how chemical bonding can affect whether a substance dissolves in a given liquid.
- CHEM.B.1.4.2 Utilize Lewis dot structures to predict the structure and bonding in simple compounds.
- CHEM.B.1.4.2 Utilize Lewis dot structures to predict the structure and bonding in simple compounds.
- BIO.A.2.2.1 Explain how carbon is uniquely suited to form biological macromolecules.
Key Terms and Definitions Students Will Learn
1. covalent bond: a bond formed when atoms share one or more pairs of electrons
2. dipole: a molecule or a part of a molecule that contains both partially positively and partially negatively charged regions
3. double covalent bond: a covalent bond in which two atoms share two pairs of electrons
4. dipole moment: (µ) the measure of net molecular polarity, which is the magnitude of the charge Q at either end of the molecular dipole times the distance r between the charges. µ=Q×r; Dipole moments tell us about the charge separation in a molecule. The larger the difference in electronegativities of bonded atoms, the larger the dipole moment. For example, NaCl has the highest dipole moment because it has an ionic bond (i.e. highest charge separation).
5. electronegativity: a measure of the ability of an atom in a chemical compound to attract electrons
6. nonpolar covalent bond: a covalent bond in which the bonding electrons are equally attracted to both bonded atoms
7. polar covalent bond: a covalent bond in which a pair of electrons are shared unequally by two atoms resulting in a partially charged molecule.
8. Non polar covalent bond: a covalent bond in which a pair of electrons are shared evenly by to atoms resulting in no net charge.
9. periodicity: synonym for periodic trends; recurring trends that are seen in the element properties; The periodicity of these properties follows trends as you move across a row or period of the periodic table or down a column or group: Moving Left → Right Ionization Energy Increases Electronegativity Increases Atomic Radius Decreases Moving Top → Bottom Ionization Energy Decreases Electronegativity Decreases Atomic Radius Increases 10. single bond: a covalent bond in which two atoms share one pair of electrons
11. Double bond: a covalent bond in which two atoms share two pair of electrons
12. triple bond: a covalent bond in which two atoms share three pairs of electrons
13. unshared pair: a nonbonding pair of electrons in the valence shell of an atom; also called lone pair
14. VSEPR theory: a theory that predicts some molecular shapes based on the idea that pairs of valence electrons surrounding an atom repel each other
15. molecular modelling: a synonym for VSEPR theory, a practice of science in that we build models to understand phenomenon
16. Molecular orbital: The region of high probability that is occupied by an individual electron as it travels with a wavelike motion in the three-dimensional space around one of two or more associated nuclei
17. Bond length: the distance between two bonded atoms at their minimum potential energy; the average distance between the nuclei of two bonded atoms.
18. Bond energy: The energy required to break the bonds in 1 mole of a chemical compound
19. Valence electron: An electron that is found in the outermost shell of an atom and determines chemical properties
20. Lewis structure: A structural formula in which electrons are represented by dots; dot pairs or dashes between two atomic symbols represent pairs in covalent bonds.
21. Resonance structure: Any one of two or more possible valence shell electron configurations of the same compound that have identical geometry but different arrangements of electrons.
22. Tetrahedral: molecular shape containing a central atom bonded to four other atoms. The angle between the bonded atoms relative to the central atom is ~109° in any direction.
Chemistry: Term 2 Unit 6
Topic: Mole and Chemical Composition
- CHEM.B.1.1.1 Apply the mole concept to representative particles (e.g., counting, determining mass of atoms, ions, molecules, and/or formula units).
- CHEM.B.1.2.1 Determine the empirical and molecular formulas of compounds.
- CHEM.B.1.2.2 Apply the law of definite proportions to the classification of elements and compounds as pure substances.
- CHEM.B.1.2.3 Relate the percent composition and mass of each element present in a compound.
Key Terms and Definitions Students Will Learn
1. avogadro’s number: 6.022 × 1023, the number of atoms or molecules in 1 mol
2. empirical formula: a chemical formula that shows the composition of a compound in terms of the relative numbers and kinds of atoms in the simplest ratio
3. excess reactant: the substance that is not used up completely in a reaction
4. limiting reactant: the substance that controls the quantity of product that can form in a chemical reaction
5. molar mass: the mass in grams of 1 mol of a substance
6. mole: the SI base unit used to measure the amount of a substance whose number of particles is the same as the number of atoms of carbon in exactly 12 g of carbon-12
7. molecular formula: a chemical formula that shows the number and kinds of atoms in a molecule, but not the arrangement of the atoms
8. Isotope: atoms of the same element type having different number of protons.
9. Average atomic mass: the percentage weight average atomic mass of the different isotope.
10. percent composition: the percentage by mass of each element in a compound.
Chemistry: Term 2 Unit 7
Topic: Chemical Reactions- Chemical Transformations
- CHEM.A.1.1.1 Classify physical or chemical changes within a system in terms of matter and/or energy.
- CHEM.A.1.1.4 Relate the physical properties of matter to its atomic or molecular structure.
- CHEM.A.1.1.5 Apply a systematic set of rules (IUPAC) for naming compounds and writing chemical formulas (e.g., binary covalent, binary ionic, ionic compounds containing polyatomic ions).
- CHEM.A.2.2.2 Predict characteristics of an atom or an ion based on its location on the periodic table (e.g., number of valence electrons, potential types of bonds, reactivity).
- CHEM.B.1.3.1 Explain how atoms combine to form compounds through ionic and covalent bonding.
- CHEM.B.1.3.2 Classify a bond as being polar covalent, non-polar covalent, or ionic.
- CHEM.B.1.3.3 Use illustrations to predict the polarity of a molecule.
- CHEM.B.1.4.1 Recognize and describe different types of models that can be used to illustrate the bonds that hold atoms together in a compound (e.g., computer models, ball-and-stick models, graphical models, solid-sphere models, structural formulas, skeletal formulas, Lewis dot structures).
- CHEM.B.1.4.2 Utilize Lewis dot structures to predict the structure and bonding in simple compounds.
- CHEM.B.2.1.3 Classify reactions as synthesis, decomposition, single replacement, double replacement, or combustion.
- CHEM.B.2.1.4 Predict products of simple chemical reactions (e.g., synthesis, decomposition, single replacement, double replacement, combustion).
- CHEM.B.2.1.5 Balance chemical equations by applying the Law of Conservation of Matter.
Key Terms and Definitions Students Will Learn
1. chemical equation: a representation of a chemical reaction that uses symbols to show the relationship between the reactants and the products
2. chemical reaction: the process by which one or more substances change to produce one or more different substances
3. coefficient: a small whole number that appears as a factor in front of a formula in a chemical equation
4. combustion reaction: the oxidation reaction of an organic compound, in which heat is released
5. decomposition reaction: a reaction in which a single compound breaks down to form two or more simpler substances
6. double replacement reaction: a reaction in which a gas, a solid precipitate, or a molecular compound forms from the apparent exchange of atoms or ions between two compounds
7. endothermic: describes a process in which heat is absorbed from the environment
8. exothermic: describes a process in which a system releases heat into the environment
9. product: a substance that forms in a chemical reaction
10. precipitate: an insoluble solid formed in a chemical reaction
11. reactant: a substance or molecule that participates in a chemical reaction
12. single replacement reaction: a reaction in which a single element reacts with a compound and displaces another element from the compound
13. synthesis reaction: a reaction in which two or more substances combine to form a new compound
14. transformation arrow: indicates the flow of the reaction , sometime referred to as yield arrow but indicates transformations that occur either forward or in reverse
15. net ionic equation: chemical equation for a reaction which lists only those species participating in the reaction.
16. spectator ions: show only those chemical species participating in a chemical reaction; spectator ions are present on the reactant side and on the product side in exactly the same form
17. nuclear fission: the splitting of the nucleus of a large atom into two or more fragments; releases additional neutrons and energy
18. nuclear fusion: the combination of the nuclei of small atoms to form a larger nucleus; releases energy
19. half life: the time required for half of a sample of a radioactive substance to disintegrate by radioactive decay or by natural processes
20. isotope: an atom that has the same number of protons (atomic number) as other atoms of the same element do but that has a different number of neutrons (atomic mass)
21. radioactivity: the process by which an unstable nucleus emits one or more particles or energy in the form of electromagnetic radiation
22. radiation: the emission of energy as electromagnetic waves or as moving subatomic particles, especially high-energy particles that cause ionization
Chemistry: Term 3 Unit 8
Topic: Stoichiometry
- CHEM.A.1.1.5 Apply a systematic set of rules (IUPAC) for naming compounds and writing chemical formulas (e.g., binary covalent, binary ionic, ionic compounds containing polyatomic ions).
- CHEM.B.1.1.1 Apply the mole concept to representative particles (e.g., counting, determining mass of atoms, ions, molecules, and/or formula units).
- CHEM.B.1.2.1 Determine the empirical and molecular formulas of compounds.
- CHEM.B.1.2.2 Apply the law of definite proportions to the classification of elements and compounds as pure substances.
- CHEM.B.1.2.3 Relate the percent composition and mass of each element present in a compound.
- CHEM.B.2.1.1 Describe the roles of limiting and excess reactants in chemical reactions.
- CHEM.B.2.1.2 Use stoichiometric relationships to calculate the amounts of reactants and products involved in a chemical reaction
Key Terms and Definitions Students Will Learn
1. avogadro’s number: 6.02 × 1023, the number of atoms or molecules in 1 mol
2. empirical formula: a chemical formula that shows the composition of a compound in terms of the relative numbers and kinds of atoms in the simplest ratio
3. excess reactant: the substance that is not used up completely in a reaction
4. limiting reactant: the substance that controls the quantity of product that can form in a chemical reaction
5. molar mass: the mass in grams of 1 mol of a substance
6. mole: the SI base unit used to measure the amount of a substance whose number of particles is the same as the number of atoms of carbon in exactly 12 g of carbon-12
7. molecular formula: a chemical formula that shows the number and kinds of atoms in a molecule, but not the arrangement of the atoms
8. percent composition: the percentage by mass of each element in a compound
9. mole ratio: use the coefficients within the balanced equation to setup the mole ratio. The mole ratio is the ratio of the unknown substance’s coefficient over the known substance’s coefficient.
Chemistry: Term 3 Unit 9
Topic: Causes of Change & Intermolecular Forces
- CHEM.A.1.1.1 Classify physical or chemical changes within a system in terms of matter and/or energy.
Key Terms and Definitions Students Will Learn
1. Heat: the energy transferred between objects that are at different temperatures.
2. Temperature: a measure of how hot (or cold) something is; specifically, a measure of the average kinetic energy of the particles in an object.
3. Enthalpy: the sum of the internal energy of a system plus the product of the system’s volume multiplied by the pressure that the system exerts on its surroundings.
4. Thermodynamics: the branch of science concerned with the energy changes that accompany chemical and physical changes.
5. Calorimetry: the measurement of heat-related constants, such as specific heat or latent heat.
6. Calorimeter: a device used to measure the heat absorbed or released in a chemical or physical change.
7. Hess’s law: the law that states the the amount of heat released or absorbed in a chemical reaction does not depend on the number of steps of the reaction; or the total energy of a reaction is the algebraic sum of the reaction parts.
8. Entropy: a measure of the randomness or disorder of a system.
9. Gibbs energy: the energy in a system that is available for work; used to determine the spontaneity of a reaction.
10. Free energy: another name for Gibbs energy; or Gibbs free energy.
11. Surface tension: the force that acts on the surface of a liquid and that tends to minimize the area of the surface.
12. Sublimation: the process in which a solid changes directly into a gas.
13. Deposition: the process in which a gas changes directly into a solid.
14. Intermolecular forces: the forces of attraction between molecules
15. Dipole - Dipole forces: interactions between polar molecules.
16. Hydrogen bond: the intermolecular force occurring when a hydrogen atom, being covalently bonded to a highly electronegative element, is attracted to a lone electron pair of another molecule. The hydrogen bond is a special type and strongest of the dipoles. a. H-bond: abbreviation for hydrogen bond.
17. London dispersion force: the intermolecular attraction resulting from the uneven distribution of electrons and the creation of temporary dipoles. a. Induced dipole: another name for the London dispersion force.
18. Phase: a part of matter that is uniform, i.e. as a solid, liquid, or gas.
19. Equilibrium: the state in which a chemical process and the reverse chemical process occur at the same rate such that the concentrations of reactants and products do not change.
20. Vapor pressure: the partial pressure exerted by a vapor that is in equilibrium with its liquid state at a given temperature.
21. Phase diagram: a graph of the relationship between the physical state of a substance and the temperature and pressure of the substance.
22. Triple point: the temperature and pressure conditions at which the solid, liquid, and gaseous phases of a substance coexist at equilibrium.
23. Critical point: the pressure at which the gas and liquid states of a substance become identical and form one phase called a supercritical fluid.
Chemistry: Term 3 Unit 10
Topic: Gases Part 1
- CHEM.B.2.2.1 Utilize mathematical relationships to predict changes in the number of particles, the temperature, the pressure, and the volume in a gaseous system (i.e., Boyle’s law, Charles’s law, Dalton’s law of partial pressures, the combined gas law, and the ideal gas law).
- CHEM.B.2.2.2 Predict the amounts of reactants and products involved in a chemical reaction using molar volume of a gas at STP.
Key Terms and Definitions Students Will Learn
1. diffusion - the movement of particles from regions of higher density to regions of lower density
2. effusion - the passage of a gas under pressure through a tiny opening
3. fluid - a substance that has no fixed shape and yields easily to external pressure; a gas or (especially) a liquid
4. ideal gas - an imaginary gas whose particles are infinitely small and do not interact with each other
5. Kelvins - the SI base unit of thermodynamic temperature - (not degrees Kelvin) the unit for temperature that is required for all the ideal gas law calculations; Kelvin = K + 273.
6. kinetic energy - the energy of an object that is due to the object’s motion
7. molar volume - the volume occupied by one mole of a substance at a given temperature and pressure
8. partial pressure - the pressure of each gas in a mixture
9. pascal - the SI unit of pressure; equal to the force of 1 N exerted over an area of 1 m2 (abbreviation, Pa)
10. pressure – the amount of force exerted per unit area
11. Newton – the SI unit for force that will increase the speed of a 1 kg mass by 1 m/s.
12. Standard Temperature and Pressure – STP is 0oC or 273K and 1 atm.
13. Ideal Gas Law – PV=nRT - no gas perfectly obeys all 4 gas laws under all conditions so scientists constructed a theoretical ‘ideal gas’ with several assumptions. This law works well as a predictor of behavior for most gases and most conditions. An idea gas, unlike a real gas, does not condense to a liquid at low temperatures, does not have forces of attraction or repulsion between particles, and is composed of particles that have no volume
14. Ideal Gas Law Constant – R 8.314 liter kPascal / mole Kelvin.
15. Gas Stoichiometry – conversion of grams to liters of a gas using molar relationships
Chemistry: Term 4 Unit 10
Topic: Gases Part 2
- CHEM.B.2.2.1 Utilize mathematical relationships to predict changes in the number of particles, the temperature, the pressure, and the volume in a gaseous system (i.e., Boyle’s law, Charles’s law, Dalton’s law of partial pressures, the combined gas law, and the ideal gas law).
- CHEM.B.2.2.2 Predict the amounts of reactants and products involved in a chemical reaction using molar volume of a gas at STP.
Key Terms and Definitions Students Will Learn
1. diffusion - the movement of particles from regions of higher density to regions of lower density
2. effusion - the passage of a gas under pressure through a tiny opening
3. fluid - a substance that has no fixed shape and yields easily to external pressure; a gas or (especially) a liquid
4. ideal gas - an imaginary gas whose particles are infinitely small and do not interact with each other
5. Kelvins - the SI base unit of thermodynamic temperature - (not degrees Kelvin) the unit for temperature that is required for all the ideal gas law calculations; Kelvin = K + 273.
6. kinetic energy - the energy of an object that is due to the object’s motion
7. molar volume - the volume occupied by one mole of a substance at a given temperature and pressure
8. partial pressure - the pressure of each gas in a mixture
9. pascal - the SI unit of pressure; equal to the force of 1 N exerted over an area of 1 m2 (abbreviation, Pa)
10. pressure – the amount of force exerted per unit area
11. Newton – the SI unit for force that will increase the speed of a 1 kg mass by 1 m/s.
12. Standard Temperature and Pressure – STP is 0oC or 273K and 1 atm.
13. Ideal Gas Law – PV=nRT - no gas perfectly obeys all 4 gas laws under all conditions so scientists constructed a theoretical ‘ideal gas’ with several assumptions. This law works well as a predictor of behavior for most gases and most conditions. An idea gas, unlike a real gas, does not condense to a liquid at low temperatures, does not have forces of attraction or repulsion between particles, and is composed of particles that have no volume
14. Ideal Gas Law Constant – R 8.314 liter kPascal / mole Kelvin.
15. Gas Stoichiometry – conversion of grams to liters of a gas using molar relationships
Chemistry: Term 4 Unit 11
Topic: Properties of Solutions
- CHEM.A.1.1.3 Utilize significant figures to communicate the uncertainty in a quantitative observation.
- CHEM.A.1.2.1 Compare properties of solutions containing ionic or molecular solutes (e.g., dissolving, dissociating).
- CHEM.A.1.2.2 Differentiate between homogeneous and heterogeneous mixtures (e.g., how such mixtures can be separated).
- CHEM.A.1.2.3 Describe how factors (e.g., temperature, concentration, surface area) can affect solubility.
- CHEM.A.1.2.4 Describe various ways that concentration can be expressed and calculated (e.g., molarity, percent by mass, percent by volume).
- CHEM.A.1.2.5 Describe how chemical bonding can affect whether a substance dissolves in a given liquid.
Key Terms and Definitions Students Will Learn
1. colligative property- a property that is determined by the number of particles present in a system but that is independent of the properties of the particles themselves.
2. colloid- a mixture consisting of tiny particles that are intermediate in size between those in solutions and those in suspensions and that are suspended in a liquid, solid, or gas.
3. concentration- the amount of a particular substance in a given quantity of a mixture, solution, or gas.
4. conductivity- the ability to conduct an electric current
5. dissociation- the separating of a molecule into simpler molecules, atoms, radicals, or ions
6. electrolyte- a substance that dissolves in water to give a solution that conducts an electric current.
7. molarity- a concentration unit of a solution expressed as moles of solute dissolved per liter of solution.
8. nonelectrolyte- a liquid or solid substance that does not allow an electric current.
9. saturated solution- a solution that cannot dissolve anymore solute under the given conditions
10. solution- a homogeneous mixture of two or more substances uniformly dispersed through-out a single phase.
11. solute- in a solution, the substance which the solvent dissolves.
12. solvent- in a solution, the substance in which the solute dissolves.
13. solubility- the ability of one substance to dissolve in another at a given temperature and pressure; expressed in terms of the amount of solute that will dissolve in a given amount of of solvent to produce a saturated solution.
14. suspension- a mixture in which particles of a material are more or less evenly dispersed throughout a liquid or gas.
15. unsaturated solution- a solution that contains less solute than a saturated solution does and that is able to dissolve an additional solute.
16. concentration - amount of particular substance in a given quantity of solution
17. supersaturated solution - a solution holding more dissolved solute than what is required to reach equilibrium at a given temperature
18. miscible - liquids completely soluble with eachother
19. immiscible - 2 or more liquids that do not mix
20. solubility equilibrium - physical state in which the opposing processes of dissolution and crystallization of a solute occur at equal rates
21. surfactant - compound that concentrates at the boundary surface between 2 immiscible phases, solid-liquid, liquid-liquid,or liquid-gas (detergent - water soluble surfactant, soap - a surfactant and type of detergent)
22. emulsion - mixture of 2 or more immiscible liquids in which one liquid is dispersed
Chemistry: Term 4 Unit 12
Topic: Equilibrium and Chemical Kinetics
- BIO.A.4.1.2 Compare and contrast the mechanisms that transport materials across the plasma membrane (i.e., passive transport -- diffusion, osmosis, facilitated diffusion; active transport -- pumps, endocytosis, exocytosis).
- CHEM.A.1.2.1 Compare properties of solutions containing ionic or molecular solutes (e.g., dissolving, dissociating).
- CHEM.A.1.2.2 Differentiate between homogeneous and heterogeneous mixtures (e.g., how such mixtures can be separated).
- CHEM.A.1.2.4 Describe various ways that concentration can be expressed and calculated (e.g., molarity, percent by mass, percent by volume).
- CHEM.B.2.1.2: Use stoichiometric relationships to calculate the amounts of reactants and products involved in a chemical reaction.
Key Terms and Definitions Students Will Learn
1. activation energy - minimum energy required to start a chemical reaction
2. endothermic - heat absorbed from environment
3. exothermic - system releases heat
4. activated complex - transition state of a reaction;
5. reaction rate - rate at which chemical reaction occurs; measured by rate of formation of the products or the rate of disappearance of the reactants
6. Le Chatelier’s principle - a system in equilibrium will oppose a change in a way that helps eliminate the change
7. reversible reaction - chemical reaction in which the products reform the original reactants
8. chemical equilibrium - state of balance in which the rate of a forward reaction equals the rate of the reverse reaction and the concentrations of products and reactants remain unchanged
9. equilibrium constant - number that relates the concentrations of starting materials and products of a reversible chemical reaction to one another at a given temperature
10. catalyst - substance that changes the rate of a chemical reaction without being consumed or changed significantly
11. common ion effect - addition of an ion common to 2 solutes brings about precipitation or reduces ionization
Chemistry: Term 4 Unit 13
Topic: Acids and Bases
- CHEM.A.1.2.3 Describe how factors (e.g., temperature, concentration, surface area) can affect solubility.
- CHEM.A.1.2.1 Compare properties of solutions containing ionic or molecular solutes (e.g., dissolving, dissociating).
- CHEM.A.1.2.4 Describe various ways that concentration can be expressed and calculated (e.g., molarity, percent by mass, percent by volume).
Key Terms and Definitions Students Will Learn
1. strong acid - acid that ionizes completely in a solvent
2. weak acid - acid that releases few hydrogen ions in aqueos solution
3. strong base - a base that ionizes completely in a solvent
4. weak base - releases few hydroxidd ions in aqueous solution
5. Bronsted-Lowry acid/base - donates a proton
6. conjugate acid/base - acid forms when a base gains a proton; base forms when an acid loses a proton
7. neutralization reaction - A chemical reaction in which an acid and base react to generate water and a salt
8. pH - Potential of Hydrogen (pH) - The quantity of hydrogen or hydroxyl ions in a solution determines whether the solution is acid or alkaline. Using a logarithmic scale, pH measures, the relative alkalinity or acidity of a solution. The pH scale was first proposed by S.P.L Sorensen, a Dane, in 1909 to identify the concentration of hydrogen ions (H+) and hydroxide ions (OH -). Of the several methods used to measure pH, two of the simplest are litmus paper and liquid acid-base indicators. Two types of litmus paper, red and blue, are coated with a dye that changes color in the presence of acids and bases. Red litmus paper will turn blue in the presence of a base. Blue litmus paper will turn red in the presence of an acid. http://www.waterqualityplus.com/ph.htm
9. equivalence point - point at which the 2 solutions used in a titration are present in chemically equivalent amounts
10. titration - method to determine the concentration of a substance in a solution by adding a solution of known volume and concentration until the reaction is completed, which is usually indicated by a color change
Does Pennsylvania Award Credit for Passing the AP Chemistry Exam?
Pennsylvania offers high school students AP Chemistry courses and students may be eligible to earn both High School and College credits for AP Chemistry courses. With a score of 3 or above on the AP Chemistry exam, students may be eligible to apply credits as equivalent to state Algebra and Biology requirements. For specific information, it is best to contact the school directly.
Does Pennsylvania Award Credit for Passing the IB Chemistry Exam?
Pennsylvania offers high school students IB Chemistry courses and students may be eligible to earn both High School and College credits for IB Chemistry courses. With a score of 4 or above IB Chemistry, students may be eligible to apply credits as equivalent to state Algebra and Biology requirements. For specific information, it is best to contact the school directly.