The modern model of the atom has evolved over a long period of time through the work of many scientists.
Each atom has a nucleus, with an overall positive charge, surrounded by negatively charged electrons.
Subatomic particles contained in the nucleus include protons and neutrons.
The proton is positively charged, and the neutron has no charge. The electron is negatively charged.
Protons and electrons have equal but opposite charges. The number of protons equals the number of electrons in an atom.
The mass of each proton and each neutron is approximately equal to one atomic mass unit. An electron is much less massive than a proton or a neutron.
The number of protons in an atom (atomic number) identifies the element. The sum of the protons and neutrons in an atom (mass number) identifies an isotope. Common notations that represent isotopes include: 14 C, 14 C, carbon-14, C-14. 6
In the wave-mechanical model (electron cloud model) the electrons are in orbitals, which are defined as the regions of the most probable electron location (ground state).
Each electron in an atom has its own distinct amount of energy.
When an electron in an atom gains a specific amount of energy, the electron is at a higher energy state (excited state).
When an electron returns from a higher energy state to a lower energy state, a specific amount of energy is emitted. This emitted energy can be used to identify an element.
The outermost electrons in an atom are called the valence electrons. In general, the number of valence electrons affects the chemical properties of an element.
Atoms of an element that contain the same number of protons but a different number of neutrons are called isotopes of that element.
The average atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.
Stability of an isotope is based on the ratio of neutrons and protons in its nucleus. Although most nuclei are stable, some are unstable and spontaneously decay, emitting radiation.
Spontaneous decay can involve the release of alpha particles, beta particles, positrons, and/ or gamma radiation from the nucleus of an unstable isotope. These emissions differ in mass, charge, ionizing power, and penetrating power.
Matter is classified as a pure substance or as a mixture of substances.
A pure substance (element or compound) has a constant composition and constant properties throughout a given sample, and from sample to sample.
Mixtures are composed of two or more different substances that can be separated by physical means. When different substances are mixed together, a homogeneous or heterogeneous mixture is formed.
The proportions of components in a mixture can be varied. Each component in a mixture retains its original properties.
Elements are substances that are composed of atoms that have the same atomic number. Elements cannot be broken down by chemical change.
Elements can be classified by their properties and located on the Periodic Table as metals, nonmetals, metalloids (B, Si, Ge, As, Sb, Te), and noble gases.
Elements can be differentiated by physical properties. Physical properties of substances, such as density, conductivity, malleability, solubility, and hardness, differ among elements.
Elements can also be differentiated by chemical properties. Chemical properties describe how an element behaves during a chemical reaction.
The placement or location of an element on the Periodic Table gives an indication of the physical and chemical properties of that element. The elements on the Periodic Table are arranged in order of increasing atomic number.
For Groups 1, 2, and 13-18 on the Periodic Table, elements within the same group have the same number of valence electrons (helium is an exception) and therefore similar chemical properties.
The succession of elements within the same group demonstrates characteristic trends: differences in atomic radius, ionic radius, electronegativity, first ionization energy, metallic/ nonmetallic properties.
The succession of elements across the same period demonstrates characteristic trends: differences in atomic radius, ionic radius, electronegativity, first ionization energy, metallic/ nonmetallic properties.
A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system.
Compounds can be differentiated by their physical and chemical properties.
Types of chemical formulas include empirical, molecular, and structural.
Organic compounds contain carbon atoms, which bond to one another in chains, rings, and networks to form a variety of structures. Organic compounds can be named using the IUPAC system.
Hydrocarbons are compounds that contain only carbon and hydrogen. Saturated hydrocarbons contain only single carbon-carbon bonds. Unsaturated hydrocarbons contain at least one multiple carbon-carbon bond.
Organic acids, alcohols, esters, aldehydes, ketones, ethers, halides, amines, amides, and amino acids are categories of organic compounds that differ in their structures. Functional groups impart distinctive physical and chemical properties to organic compounds.
Isomers of organic compounds have the same molecular formula, but different structures and properties.
The structure and arrangement of particles and their interactions determine the physical state of a substance at a given temperature and pressure.
The three phases of matter (solids, liquids, and gases) have different properties.
Entropy is a measure of the randomness or disorder of a system. A system with greater disorder has greater entropy.
Systems in nature tend to undergo changes toward lower energy and higher entropy.
Differences in properties such as density, particle size, molecular polarity, boiling and freezing points, and solubility permit physical separation of the components of the mixture.
A solution is a homogeneous mixture of a solute dissolved in a solvent. The solubility of a solute in a given amount of solvent is dependent on the temperature, the pressure, and the chemical natures of the solute and solvent.
The concentration of a solution may be expressed in molarity (M), percent by volume, percent by mass, or parts per million (ppm).
The addition of a nonvolatile solute to a solvent causes the boiling point of the solvent to increase and the freezing point of the solvent to decrease. The greater the concentration of solute particles, the greater the effect.
An electrolyte is a substance which, when dissolved in water, forms a solution capable of conducting an electric current. The ability of a solution to conduct an electric current depends on the concentration of ions.
The acidity or alkalinity of an aqueous solution can be measured by its pH value. The relative level of acidity or alkalinity of these solutions can be shown by using indicators.
On the pH scale, each decrease of one unit of pH represents a tenfold increase in hydronium ion concentration.
Behavior of many acids and bases can be explained by the Arrhenius theory. Arrhenius acids and bases are electrolytes.
Arrhenius acids yield H + (aq), hydrogen ion as the only positive ion in an aqueous solution. The hydrogen ion may also be written as H 3 O + (aq), hydronium ion.
Arrhenius bases yield OH -(aq), hydroxide ion as the only negative ion in an aqueous solution.
In the process of neutralization, an Arrhenius acid and an Arrhenius base react to form a salt and water.
There are alternate acid-base theories. One theory states that an acid is an H + donor and a base is an H + acceptor.
Titration is a laboratory process in which a volume of a solution of known concentration is used to determine the concentration of another solution.
A physical change results in the rearrangement of existing particles in a substance. A chemical change results in the formation of different substances with changed properties.
Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement.
Types of organic reactions include addition, substitution, polymerization, esterification, fermentation, saponification, and combustion.
An oxidation-reduction (redox) reaction involves the transfer of electrons (e -).
Reduction is the gain of electrons.
A half-reaction can be written to represent reduction.
Oxidation is the loss of electrons.
A half-reaction can be written to represent oxidation.
Oxidation numbers (states) can be assigned to atoms and ions. Changes in oxidation numbers indicate that oxidation and reduction have occurred.
An electrochemical cell can be either voltaic or electrolytic. In an electrochemical cell, oxidation occurs at the anode and reduction at the cathode.
A voltaic cell spontaneously converts chemical energy to electrical energy.
An electrolytic cell requires electrical energy to produce a chemical change. This process is known as electrolysis.
In all chemical reactions there is a conservation of mass, energy, and charge.
In a redox reaction the number of electrons lost is equal to the number of electrons gained.
A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction.
The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound.
The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram-formula mass) of a substance equals one mole of that substance.
The percent composition by mass of each element in a compound can be calculated mathematically.
The concept of an ideal gas is a model to explain the behavior of gases. A real gas is most like an ideal gas when the real gas is at low pressure and high temperature.
are in random, constant, straight-line motion.
are separated by great distances relative to their size; the volume of the gas particles is considered negligible.
have no attractive forces between them.
have collisions that may result in a transfer of energy between gas particles, but the total energy of the system remains constant.
Kinetic molecular theory describes the relationships of pressure, volume, temperature, velocity, and frequency and force of collisions among gas molecules.
Collision theory states that a reaction is most likely to occur if reactant particles collide with the proper energy and orientation.
Equal volumes of gases at the same temperature and pressure contain an equal number of particles.
The rate of a chemical reaction depends on several factors: temperature, concentration, nature of the reactants, surface area, and the presence of a catalyst.
A catalyst provides an alternate reaction pathway, which has a lower activation energy than an uncatalyzed reaction.
Some chemical and physical changes can reach equilibrium.
At equilibrium the rate of the forward reaction equals the rate of the reverse reaction. The measurable quantities of reactants and products remain constant at equilibrium.
LeChatelier's principle can be used to predict the effect of stress (change in pressure, volume, concentration, and temperature) on a system at equilibrium.
Energy can exist in different forms, such as chemical, electrical, electromagnetic, thermal, mechanical, nuclear.
Chemical and physical changes can be exothermic or endothermic.
Energy released or absorbed during a chemical reaction can be represented by a potential energy diagram.
Energy released or absorbed during a chemical reaction (heat of reaction) is equal to the difference between the potential energy of the products and potential energy of the reactants.
Heat is a transfer of energy (usually thermal energy) from a body of higher temperature to a body of lower temperature. Thermal energy is the energy associated with the random motion of atoms and molecules.
Temperature is a measurement of the average kinetic energy of the particles in a sample of material. Temperature is not a form of energy.
The concepts of kinetic and potential energy can be used to explain physical processes that include: fusion (melting), solidification (freezing), vaporization (boiling, evaporation), condensation, sublimation, and deposition.
use models to describe the structure of an atom
relate experimental evidence (given in the introduction of Key Idea 3) to models of the atom
determine the number of protons or electrons in an atom or ion when given one of these values
calculate the mass of an atom, the number of neutrons or the number of protons, given the other two values
distinguish between ground state and excited state electron configurations, e.g., 2-8-2 vs. 2-7-3
identify an element by comparing its bright-line spectrum to given spectra
distinguish between valence and non-valence electrons, given an electron configuration, e.g., 2-8-2
draw a Lewis electron-dot structure of an atom
determine decay mode and write nuclear equations showing alpha and beta decay
interpret and write isotopic notation
given an atomic mass, determine the most abundant isotope
calculate the atomic mass of an element, given the masses and ratios of naturally occurring isotopes
classify elements as metals, nonmetals, metalloids, or noble gases by their properties
compare and contrast properties of elements within a group or a period for Groups 1, 2, 13-18 on the Periodic Table
determine the group of an element, given the chemical formula of a compound, e.g., XCl or XCl2
explain the placement of an unknown element on the Periodic Table based on its properties
classify an organic compound based on its structural or condensed structural formula (i.e., CH3COOH or -C-C-OH)
describe the states of the elements at STP
distinguish among ionic, molecular, and metallic substances, given their properties
draw a structural formula with the functional group(s) on a straight chain hydrocarbon backbone, when given the IUPAC name for the compound
draw structural formulas for alkanes, alkenes, and alkynes containing a maximum of ten carbon atoms
use a simple particle model to differentiate among properties of solids, liquids, and gases
compare the entropy of phases of matter
describe the processes and uses of filtration, distillation, and chromatography in the separation of a mixture
interpret and construct solubility curves
apply the adage "like dissolves like" to real-world situations
interpret solution concentration data
use solubility curves to distinguish among saturated, supersaturated, and unsaturated solutions
calculate solution concentration in molarity (M), percent mass, and parts per million (ppm)
describe the preparation of a solution, given the molarity
given properties, identify substances as Arrhenius acids or Arrhenius bases
identify solutions as acid, base, or neutral based upon the pH
interpret changes in acid-base indicator color
write simple neutralization reactions when given the reactants
calculate the concentration or volume of a solution, using titration data
use particle models/diagrams to differentiate among elements, compounds, and mixtures
distinguish between chemical and physical changes
identify types of chemical reactions
determine a missing reactant or product in a balanced equation
identify organic reactions
balance equations, given the formulas of reactants and products
write and balance half-reactions for oxidation and reduction of free elements and their monatomic ions
identify and label the parts of a voltaic cell (cathode, anode, salt bridge) and direction of electron flow, given the reaction equation
identify and label the parts of an electrolytic cell (cathode, anode) and direction of electron flow, given the reaction equation
compare and contrast voltaic and electrolytic cells
use an activity series to determine whether a redox reaction is spontaneous
balance equations, given the formulas for reactants and products
interpret balanced chemical equations in terms of conservation of matter and energy
create and use models of particles to demonstrate balanced equations
calculate simple mole-mole stoichiometry problems, given a balanced equation
determine the empirical formula from a molecular formula
determine the mass of a given number of moles of a substance
determine the molecular formula, given the empirical formula and the molecular mass
calculate the formula mass and gram-formula mass
determine the number of moles of a substance, given its mass