Ohio Requirements for Passing High School Chemistry | General Chemistry 1

Is Chemistry Required in High School in Ohio?

Students must take and earn a state minimum of 20 credits in specific subjects, with a minimum of 3 Science credits. Science units must include one unit of physical sciences, one unit of life sciences, and one unit of advanced study in one or more of the following sciences: chemistry, physics or other physical science; advanced biology or other life science; astronomy, physical geology or other earth or space science. Students can meet one of the following options to fulfill high school graduation requirements:

 

OPTION 1
Satisfy one of the three original pathways to graduation that were in place when you entered high school. The pathways include:  

  • Ohio’s State Tests - Earn at least 18 points on seven end-of-course state tests. Each test score earns you up to five graduation points. You must have a minimum of four points in math, four points in English, and six points across science and social studies.
  • Industry credential and workforce readiness - Earn a minimum of 12 points by receiving a State Board of Education-approved, industry-recognized credential or group of credentials in a single career field and earn the required score on WorkKeys, a work-readiness test. The state of Ohio will pay one time for you to take the WorkKeys test.  
  • College and career readiness tests - Earn remediation-free scores in mathematics and English language arts on either the ACT or SAT.  

 

OPTION 2

Satisfy the new graduation requirements for the classes of 2023 and beyond by: 

The structure of Ohio’s Learning Standards for Science is significantly different from NGSS, but the research that provided the framework from which each was developed is the same. Both sets of standards address similar science content, skills, and ways of thinking. Students who study chemistry topics in their core high school classes will learn about the following chemistry subjects:

 

Ohio High School Chemistry Course Description

Chemistry is a high school level course, which satisfies the Ohio Core science graduation requirements of Ohio Revised Code Section 3313.603. This section of Ohio law requires three units of science. Each course should include inquiry-based laboratory experience that engages students in asking valid scientific questions and gathering and analyzing information.

This course introduces students to key concepts and theories that provide a foundation for further study in other sciences as well as advanced science disciplines. Chemistry comprises a systematic study of the predictive physical interactions of matter and subsequent events that occur in the natural world. The study of matter through the exploration of classification, its structure and its interactions is how this course is organized.

Investigations are used to understand and explain the behavior of matter in a variety of inquiry and design scenarios that incorporate scientific reasoning, analysis, communication skills and real-world applications. An understanding of leading theories and how they have informed current knowledge prepares students with higher-order cognitive capabilities of evaluation, prediction, and application.

Course Content

The following information may be taught in any order; there is no ODE-recommended sequence.

C.PM: STRUCTURE AND PROPERTIES OF MATTER

  • C.PM.1: Atomic structure
    • Evolution of atomic models/theory
    • Electrons • Electron configurations
  • C.PM.2: Periodic Table
    • Properties
    • Trends
  • C.PM.3: Chemical bonding
    • Ionic
    • Polar/covalent
  • C.PM.4: Representing compounds
    • Formula writing
    • Nomenclature
    • Models and shapes (Lewis structures, ball and stick, molecular geometries)
  • C.PM.5: Quantifying matter
  • C.PM.6: Intermolecular forces of attraction
    • Types and strengths
    • Implications for properties of substances
      • Melting and boiling point
      • Solubility
      • Vapor pressure

C.IM: INTERACTIONS OF MATTER

  • C.IM.1: Chemical reactions
    • Types of reactions
    • Kinetics
    • Energy
    • Equilibrium
    • Acids/bases
  • C.IM.2: Gas laws
    • Pressure, volume, and temperature
    • Ideal gas law
  • C.IM.3: Stoichiometry
    • Molecular calculations
    • Solutions
    • Limiting reagents 
C.PM: STRUCTURE AND PROPERTIES OF MATTER

C.PM.1: Atomic structure

  • Evolution of atomic models/theory
  • Electrons
  • Electron configurations

201 8 CONTENT ELABORATION: STRUCTURE AND PROPERTIES OF MATTER

 

Atomic structure Physical Science included properties and locations of protons, neutrons and electrons, atomic number, mass number, cations and anions, isotopes and the strong nuclear force which holds the nucleus together. In this course, the historical development of the atomic model and the positions of electrons are explored in greater detail. Atomic models are constructed to explain experimental evidence and make predictions. The changes in the atomic model over time exemplify how scientific knowledge changes as new evidence emerges and how technological advancements like electricity extend the boundaries of scientific knowledge. Thompson’s study of electrical discharges in cathode-ray tubes led to the discovery of the electron and the development of the plum pudding model of the atom. Rutherford’s experiment, in which he bombarded gold foil with α-particles, led to the discovery that most of the atom consists of empty space with a relatively small, positively charged nucleus. Bohr used data from atomic spectra to propose a planetary model of the atom in which electrons orbit the nucleus, like planets around the sun. Later, Schrödinger used the idea that electrons travel in waves to develop a model in which electrons travel randomly in regions of space called orbitals (quantum mechanical model).

Based on the quantum mechanical model, it is not possible to predict exactly where electrons are located but there is a region of space surrounding the nucleus in which there is a high probability of finding an electron (electron cloud or orbital). Data from atomic spectra (emission and absorption) gives evidence that electrons can only exist at certain discrete energy levels and not at energies between these levels.

Atoms are usually in the ground state where the electrons occupy orbitals with the lowest available energy. However, the atom can become excited when the electrons absorb a photon with the precise amount of energy (indicated by the frequency of the photon) to move to an orbital with higher energy. Any photon without this precise amount of energy will be ignored by the electron. The atom exists in the excited state for a very short amount of time. When an electron drops back down to the lower energy level, it emits a photon that has energy equal to the energy difference between the levels. The amount of energy is indicated by the frequency of the light that is given off and can be measured. Each element has a unique emission and absorption spectrum due to its unique electron configuration and specific electron energy jumps that are possible for that element.

Being aware of the quantum mechanical model as the currently accepted model for the atom is important for science literacy as it explains and predicts subatomic interactions, but details should be reserved for more advanced study.

Electron energy levels consist of sublevels (s, p, d, and f), each with a characteristic number and shape of orbitals. Orbital diagrams and electron configuration can be constructed to show the location of the electrons in an atom using established rules. Valence electrons are responsible for most of the chemical properties of elements. In this course, electron configuration (extended and noble gas notation) and orbital diagrams can be shown for any element in the first three periods.

Although the quantum mechanical model of the atom explains the most experimental evidence, other models can still be helpful. Thinking of atoms as indivisible spheres is useful in explaining many physical properties of substances, such as the state (solid, liquid, or gas) of a substance at room temperature. Bohr’s planetary model is useful to explain and predict periodic trends in the properties of elements. 

 

C.PM.2: Periodic table

  • Properties
  • Trends

CONTENT ELABORATION: STRUCTURE AND PROPERTIES OF MATTER:

In the Physical Science course, the concept that elements are placed in order of increasing atomic number in the periodic table such that elements with similar properties are placed in the same column is introduced. How the periodic table is divided into groups, families, periods, metals, nonmetals, and metalloids is also included and will be revisited here. In this course, with more information about the electron configuration of elements, similarities in the configuration of the valence electrons for a particular group can be observed.

The electron configuration of an atom can be determined from the position on the periodic table. The repeating pattern in the electron configuration for elements on the periodic table explains many of the trends in the properties observed. Atomic theory is used to describe and explain trends in properties across periods or down columns including atomic radii, ionic radii, first ionization energies, electronegativities, and whether the element is a solid or gas at room temperature.

Additional ionization energies, electron affinities, and periodic properties of the transition elements, and the lanthanide and actinide series are reserved for more advanced study. 

 

C.PM.3: Chemical bonding

  • Ionic
  • Polar/covalent

201 8 CONTENT ELABORATION: STRUCTURE AND PROPERTIES OF MATTER

Chemical bonding Content in the Physical Science course included recognizing that atoms with unpaired electrons tend to form ionic and covalent bonds with other atoms, forming molecules, ionic lattices, or network covalent structures. In this course, electron configuration, electronegativity values, and energy considerations will be applied to bonding and the properties of materials with different types of bonding.

Atoms of many elements are more stable when they are bonded to other atoms. In such cases, as atoms bond, energy is released to the surroundings, resulting in a system with lower energy. An atom’s electron configuration, particularly the valence electrons, determines how an atom interacts with other atoms. Molecules, ionic lattices, and network covalent structures have different, yet predictable, properties that depend on the identity of the elements and the types of bonds formed.

Differences in electronegativity values can be used to predict where a bond fits on the continuum between ionic and covalent bonds. The polarity of a bond depends on the electronegativity difference and the distance between the atoms (bond length). Polar covalent bonds are introduced as an intermediary between ionic and pure covalent bonds. The concept of metallic bonding is also introduced to explain many of the properties of metals (e.g., conductivity). Since most compounds contain multiple bonds, a substance may contain more than one type of bond. Carbon atoms can bond together and with other atoms, especially hydrogen, oxygen, nitrogen, and sulfur, to form chains, rings, and branching networks that are present in a variety of important compounds, including synthetic polymers, fossil fuels, and large molecules essential to life. Detailed study of the structure of molecules responsible for life is reserved for more advanced courses.

 

C.PM.4: Representing compounds

  • Formula writing
  • Nomenclature
  • Models and shapes (Lewis structures, ball and stick, molecular geometries)

CONTENT ELABORATION: STRUCTURE AND PROPERTIES OF MATTER

Representing compounds Using the periodic table, formulas of ionic compounds containing specific elements can be predicted. This can include ionic compounds made up of elements from groups 1, 2, 17, hydrogen, oxygen, and polyatomic ions (given the formula and charge of the polyatomic ion). Given the formula, a compound can be named using conventional systems that include Greek prefixes and Roman numerals where appropriate. Given the name of an ionic or covalent substance, formulas can be written.

Many different models can be used to represent compounds including chemical formulas, Lewis structures, and ball and stick models. These models can be used to visualize atoms and molecules and to predict the properties of substances. Each type of representation provides unique information about the compound. Different representations are better suited for particular substances. Lewis structures can be drawn to represent covalent compounds using a simple set of rules and can be combined with valence shell electron pair repulsion (VSEPR) theory to predict the three-dimensional electron pair and molecular geometry of compounds. Lewis structures and molecular geometries will only be constructed for the following combination of elements: hydrogen, carbon, nitrogen, oxygen, phosphorus, sulfur, and the halogens. Organic nomenclature is reserved for more advanced courses. 

 

C.PM.5: Quantifying matter

201 8 CONTENT ELABORATION: STRUCTURE AND PROPERTIES OF MATTER

Quantifying matter In earlier grades, properties of materials were quantified with measurements that were always associated with some error. In this course, scientific protocols for quantifying the properties of matter accurately and precisely are studied. Using the International System of Units (SI), significant digits or figures, scientific notation, error analysis, and dimensional analysis are vital to scientific communication.

There are three domains of magnitude in size and time: the macroscopic (human) domain, the cosmic domain, and the submicroscopic (atomic and subatomic) domain. Measurements in the cosmic domain and submicroscopic domains require complex instruments and/or procedures.

Matter can be quantified in a way that macroscopic properties such as mass can reflect the number of particles present. Elemental samples are a mixture of several isotopes with different masses. The atomic mass of an element is calculated given the mass and relative abundance of each isotope of the element as it exists in nature. Because the mass of an atom is very small, the mole is used to translate between the atomic and macroscopic levels. A mole is equal to the number of atoms in exactly 12 grams of the isotope carbon-12. The mass of one mole of a substance is equal to its molar mass in grams. The molar mass for a substance can be used in conjunction with Avogadro’s number and the density of a substance to convert between mass, moles, volume, and number of particles of a sample. 

 

C.PM.6: Intermolecular forces of attraction

  • Types and strengths
  • Implications for properties of substances
    • Melting and boiling point
    • Solubility
    • Vapor pressure 201 8

CONTENT ELABORATION: STRUCTURE AND PROPERTIES OF MATTER

In middle school, solids, liquids, and gases were explored in relation to the spacing of the particles, motion of the particles, and strength of attraction between the particles that make up the substance. The intermolecular forces of attraction between particles that determine whether a substance is a solid, liquid, or gas at room temperature are addressed in greater detail in this course. Intermolecular attractions are generally weak when compared to intramolecular bonds, but span a wide range of strengths.

The composition of a substance and the shape and polarity of a molecule are particularly important in determining the type and strength of bonding and intermolecular interactions. Types of intermolecular attractions include London dispersion forces (present between all molecules), dipole-dipole forces (present between polar molecules), and hydrogen bonding (a special case of dipole-dipole where hydrogen is bonded to a highly electronegative atom such as fluorine, oxygen, or nitrogen), each with its own characteristic relative strength.

The configuration of atoms in a molecule determines the strength of the forces (bonds or intermolecular forces) between the particles and therefore the physical properties (e.g., melting point, boiling point, solubility, vapor pressure) of a material. For a given substance, the average kinetic energy (temperature) needed for a change of state to occur depends upon the strength of the intermolecular forces between the particles. Therefore, the melting point and boiling point depend upon the amount of energy that is needed to overcome the attractions between the particles.

Substances that have strong intermolecular forces or are made up of threedimensional networks of ionic or covalent bonds, tend to be solids at room temperature and have high melting and boiling points. Nonpolar organic molecules are held together by weak London dispersion forces. However, substances with longer chains provide more opportunities for these attractions and tend to have higher melting and boiling points. Increased branching of organic molecules results in lower melting and boiling points due to interference with the intermolecular attractions. Substances will have a greater solubility when dissolving in a solvent with similar intermolecular forces. If the substances have different intermolecular forces, they are more likely to interact with themselves than the other substance and remain separated from each other. Water is a polar molecule and it is often used as a solvent since most ionic and polar covalent substances will dissolve in it. In order for an ionic substance to dissolve in water, the attractive forces between the ions must be overcome by the dipole-dipole interactions with the water. Dissolving of a solute in water is an example of a process that is difficult to classify as a chemical or physical change and it is not appropriate to have students classify it one way or another.

Evaporation occurs when the particles with enough kinetic energy to overcome the attractive forces separate from the rest of the sample to become a gas. The pressure of these particles is called vapor pressure. Vapor pressure increases with temperature. Particles with larger intermolecular forces have lower vapor pressures at a given temperature since the particles require more energy to overcome the attractive forces between them. Molecular substances often evaporate more due to the weak attractions between the particles and can often be detected by their odor. Ionic or network covalent substances have stronger forces and are not as likely to volatilize. These substances often have little, if any, odor. Liquids boil when their vapor pressure is equal to atmospheric pressure. In solid water, there is a network of hydrogen bonds between the particles that gives it an open structure. This is why water expands as it freezes and why solid water has a lower density than liquid water. This has important implications for life (e.g., ice floating on water acts as an insulator in bodies of water to keep the temperature of the rest of the water above freezing). 

 

C.IM.1: Chemical reactions

  • Types of reactions
  • Kinetics
  • Energy
  • Equilibrium
  • Acids/bases

CONTENT ELABORATION: INTERACTIONS OF MATTER

In the Physical Science course, coefficients were used to balance simple equations. Other representations, including Lewis structures and three-dimensional models, were also used and manipulated to demonstrate the conservation of matter in chemical reactions. In this course, more complex reactions will be studied, classified, and represented with balanced chemical equations and three-dimensional models.

Classifying reactions into types can be a helpful organizational tool for recognizing patterns of what may happen when two substances are mixed. Teachers should be aware that the common reaction classifications that are often used in high school chemistry courses may lead to misconceptions because they are not based on the actual chemistry, but on surface features that can be similar from one system to another (e.g., exchanging partners), even though the underlying chemistry is not the same. However, these classifications may be useful in making predictions about what happens when two substances are mixed.

Some general types of chemical reactions are oxidation/reduction, synthesis, decomposition, single replacement, double replacement (including precipitation reactions and some acid-base neutralizations), and combustion reactions. Some reactions can fit into more than one category. For example, a single replacement reaction can also be classified as an oxidation/reduction reaction. Identification of reactions involving oxidation and reduction as well as indicating what substance is being oxidized and what is being reduced are appropriate in this course. However, balancing complex oxidation/reduction reactions is reserved for more advanced study.

Organic molecules release energy when undergoing combustion reactions and are used to meet the energy needs of society (e.g., oil, gasoline, natural gas) and to provide the energy needs of biological organisms (e.g., cellular respiration). When a reaction between two ionic compounds in an aqueous solution results in the formation of a precipitate or molecular compound, the reaction often occurs because the new ionic or covalent bonds are stronger than the original ion-dipole interactions of the ions in solution. Laboratory experiences (3-D or virtual) with different types of chemical reactions should be provided.

Reactions occur when reacting particles collide in an appropriate orientation and with sufficient energy. The rate of a chemical reaction is the change in the amount of the reactants or products in a specific period of time. Increasing the probability or effectiveness of the collisions between the particles increases the rate of the reaction. Therefore, changing the concentration of the reactants, changing the temperature or the pressure of gaseous reactants, or using a catalyst, can change the reaction rate. Likewise, the collision theory can be applied to dissolving solids in a liquid solvent and can be used to explain why reactions are more likely to occur between reactants in the aqueous or gaseous state than between solids. The rate at which a substance dissolves should not be confused with the amount of solute that can dissolve in a given amount of solvent (solubility). Mathematical treatment of reaction rates is reserved for more advanced study. Computer simulations can help visualize reactions from the perspective of the kinetic-molecular theory.

In middle school, the differences between potential and kinetic energy and the particle nature of thermal energy were introduced. For chemical systems, potential energy is in the form of chemical energy and kinetic energy is in the form of thermal energy. The total amount of chemical energy and/or thermal energy in a system is impossible to measure. However, the energy change of a system can be calculated from measurements (mass and change in temperature) from calorimetry experiments in the laboratory. Conservation of energy is an important component of calorimetry equations. Thermal energy is the energy of a system due to the movement of its particles. The thermal energy of an object depends upon the amount of matter present (mass), temperature, and chemical composition.

Some materials require little energy to change their temperature and other materials require a great deal to change their temperature by the same amount. Specific heat is a measure of how much energy is needed to change the temperature of a specific mass of material to a specific amount. Specific heat values can be used to calculate the thermal energy change, the temperature (initial, final, or change in) or mass of a material in calorimetry. Water has a particularly high specific heat capacity, which is important in regulating Earth’s temperature.

As studied in middle school, chemical energy is the potential energy associated with chemical systems. Chemical reactions involve valence electrons forming bonds to yield more stable products with lower energies. Energy is required to break interactions and bonds between the reactant atoms and energy is released when an interaction or bond is formed between the atoms in the products. Molecules with weak bonds (e.g., ATP) are less stable and tend to react to produce more stable products, releasing energy in the process. Generally, energy is transferred out of the system (exothermic) when the products have stronger bonds than the reactants and is transferred into the system (endothermic) when the reactants have stronger bonds than the products. Predictions of the energy requirements (endothermic or exothermic) of a reaction can be made given a table of bond energies. Graphic representations can be drawn and interpreted to represent the energy changes during a reaction. The role of energy in determining the spontaneity of chemical reactions is dealt with conceptually in this course. Entropy and its influence on the spontaneity of reactions are reserved for more advanced study.

All reactions are reversible to a degree and many reactions do not proceed completely toward products but appear to stop progressing before the reactants are all used up. At this point, the amounts of the reactants and the products appear to be constant and the reaction can be said to have reached dynamic equilibrium. Dynamic equilibrium means the rate of the reverse reaction is equal to the rate of the forward reaction so there is no apparent change in the reaction.

If a chemical system at equilibrium is disturbed by a change in the conditions of the system (e.g., increase or decrease in the temperature, pressure on gaseous equilibrium systems, concentration of a reactant or product), then the equilibrium system will respond by shifting to a new equilibrium state, reducing the effect of the change (Le Chatelier’s Principle). If products are removed as they are formed during a reaction, then the equilibrium position of the system is forced to shift to favor the products. In this way, an otherwise unfavorable reaction can be made to occur. Mathematical treatment of equilibrium reactions is reserved for advanced study. Computer simulations can help visualize the progression of a reaction to dynamic equilibrium and the continuation of both the forward and reverse reactions after equilibrium has been attained.

Properties of acids and bases and the ranges of the pH scale were introduced in Physical Science. In this course, the structural features of molecules are explored to further understand acids and bases. Acids often result when hydrogen is covalently bonded to an electronegative element and is easily dissociated from the rest of the molecule to bind with water to form a hydronium ion (H3O+). The acidity of an aqueous solution can be expressed as pH, where pH can be calculated from the concentration of the hydronium ion. Bases are likely to dissociate in water to form a hydroxide ion. Acids can react with bases to form a salt and water. Such neutralization reactions can be studied quantitatively by performing titration experiments. Detailed instruction about the equilibrium of acids and bases and the concept of Brønsted-Lowry and Lewis acids and bases is not the focus at this level. 

 

C.IM.2: Gas laws

  • Pressure, volume, and temperature
  • Ideal gas law

CONTENT ELABORATION: INTERACTIONS OF MATTER

The kinetic-molecular theory can be used to explain the properties of gases (pressure, temperature, and volume) through the motion and interactions of its particles. Problems can also be solved involving the changes in temperature, pressure, volume, and amount of a gas. When two of these four are kept constant, the relationship between the other two can be quantified, described, and explained using the kinetic-molecular theory. Real-world phenomena (e.g., why tire pressure increases in hot weather, why a hot air balloon rises) can be explained using this theory. When solving gas problems, the Kelvin temperature scale must be used since only in this scale is the temperature directly proportional to the average kinetic energy. The Kelvin temperature is based on a scale that has its minimum temperature at absolute zero, a temperature at which all motion theoretically stops. Since equal volumes of gases at the same temperature and pressure contain an equal number of particles (Avogadro’s law), problems can be solved for an unchanging gaseous system using the ideal gas law (PV = nRT) where R is the ideal gas constant (e.g., represented in multiple formats, 8.31 joules/(mole·K). The focus in this course is solving problems using the gas laws and understanding their applications, rather than memorizing the specific names and formulas. Deviations from ideal gaseous behavior are reserved for more advanced study. Relationships between the volume, temperature, and pressure can be explored in the laboratory or through computer simulations or virtual experiments. 

 

C.IM.3: Stoichiometry

  • Molar calculations
  • Solutions
  • Limiting reagents

CONTENT ELABORATION: INTERACTIONS OF MATTER

A stoichiometric calculation involves the conversion from the amount of one substance in a chemical reaction to the amount of another substance. The coefficients of the balanced equation indicate the ratios of the substances involved in the reaction in terms of both particles and moles.

Once the number of moles of a substance is known, amounts can be changed to mass, volume of a gas, volume of solutions and/or number of particles. Molarity is a measure of the concentration of a solution that can be used in stoichiometric calculations. When performing a reaction in the lab, the experimental yield can be compared to the theoretical yield to calculate percent yield. The concept of limiting reagents is treated conceptually. Mathematical applications can be utilized, but it is important to address the symbolic representations as well. Molality and normality are concepts reserved for more advanced study. 

 

Does Ohio Award Credit for Passing the AP Chemistry Exam?

Students can earn high school graduation credit for College Credit Plus (CCP) courses or approved AP/IB tests, in the subject area, which will satisfy the end-of-course graduation test requirement for biology. Students will need to understand the basic concepts of chemistry as building blocks for life. 

For college credit, schools usually award credit for AP Chemistry with a score of 3 or higher on the exam. To determine if a score qualifies, contact the school directly.