chem
We use balanced equations to determine quantities of reactants and products involved in chemical reactions. The coefficients tell us the relative numbers of moles of reactants that combine and products that form.
We combine coefficients to give mole ratios (or molecule ratios)that allow us to calculate moles of one reaction species from another one that is known.
If the mass in grams is given or required, conversion between grams and moles using molar mass must be added to the calculation.
When reactants are not present in exactly the right amounts for complete reaction of all reactants, only the limiting reactant reacts completely.
Because the limiting reactant is in short supply, its amount limits the amount of product that can form. This is why we use the amount of the limiting reactant to calculate the theoretical yield—the amount of product that is predicted.
If the reaction is done in the laboratory, the amount of product actually obtained can be compared to the theoretical yield by calculating a percent yield.
Energy changes accompany chemical and physical changes.
Endothermic reactions absorb heat from the surroundings, and exothermic reactions release heat to the surroundings.
The specific heat value for a substance is the amount of heat needed to raise the temperature of 1 g of a substance by 1°C.
A substance’s specific heat value, C, allows us to relate the heat change of a substance, q, its mass, m, and temperature change, ΔT,by this equation:
q=m×C×ΔT
Calorimetry is used to determine the heat change for a system (usually an object or a chemical reaction).
Calorimetry usually involves calculating the heat change of the surroundings (qsurroundings) from its measured temperature change. The law of conservation of energy is then used to relate the heat change of the system to the heat change of the surroundings: qsystem=−qsurroundings.
Heat changes for chemical reactions are usually reported in units of kJ/mol or kJ/g of a specific reactant.
Scientists interested in light led the way to our current understanding of atoms—specifically, how electrons are arranged in atoms.
Light can be described by its wavelength, frequency, or photon energy. The light we see is a small portion of the electromagnetic spectrum, the visible region.
When an element is heated or given an electric charge, the energy absorbed by its atoms is emitted as light energy. The light an element produces in this way contains only specific colors, as shown by its line spectrum.
Bohr’s model of the hydrogen atom helped scientists understand the quantized nature of electrons—that they have specific allowed energies.
Bohr showed that the four lines in the hydrogen spectrum result from electron transitions from higher-energy orbits (n=3, 4, 5, and 6) to the n=2orbit.
Bohr’s model only explained the hydrogen atom spectrum; however, the fundamental ideas led to the modern model of the atom.
The idea that the energies of electrons are quantized is retained in the modern model of the atom, where electrons are described as being in specific principal energy levels, sublevels, and orbitals.
An orbital is a three-dimensional region in space where an electron is likely to be found.
Sublevels contain sets of orbitals with the same letter designations: s, p, d, or f.
The s sublevel contains one orbital, the p sublevel contains three orbitals, the d sublevel contains five orbitals, and the fsublevel contains seven orbitals.
Orbital diagrams show the relative energies of the orbitals, and electrons occupy those orbitals following the aufbau principle, the Pauli exclusion principle, and Hund’s rule.
Electron configurations are shorthand notations that show the distribution of electrons in atoms.
The periodic table can be divided into s, p, d, and f blocks and is a useful tool for writing electron configurations.
Valence electrons are the outer electrons that are held less tightly by the nucleus than the core electrons. Valence electrons for main-group elements are the electrons in the highest principal energy level (highest n value); they are the electrons in the last-occupied s and p sublevels, so elements in the same group have the same number of valence electrons.
Electron configurations for ions are written by adjusting the number of valence electrons.
We can explain many of the trends in chemical and physical properties of elements by looking at specific atomic properties such as ionization energy and atomic size.
Ionization energy (IE) is a measure of the amount of energy required to remove the highest-energy valence electron from an atom. There is a periodic trend for ionization energy values: They increase going up a group and from left to right within a period.
Because successive electrons are more difficult to remove, successive ionization energies increase in the orderIE1<IE2<IE3, etc.
The periodic trend for atomic size is opposite that of ionization energy: atomic size increases going down a group and from right to left within a period.
Cations are smaller than their neutral atoms, and anions are larger than their neutral atoms. Because isoelectronic ions differ only in the number of protons, size decreases as the number of protons increases.
A chemical bond is the force that holds atoms together in a molecule or compound.
Different types of chemical bonds give substances different properties.
Electronegativity is a measure of the tendency for an atom to attract electrons in a bond.
The electrons in a covalent bond are not always shared equally, in which case the covalent bonds are polar.
Page 338
Chemical bonds tend to form in conformance with the octet rule.
Atoms tend to lose, gain, or share electrons to achieve a more stable electron configuration when bonding to form a compound.
According to the octet rule, atoms tend to be associated with eight valence electrons.
To obtain the noble-gas configuration of helium, hydrogen usually has two electrons when bonded. There are some other exceptions to the octet rule.
Ionic bonds form through transfer of electrons.
Ionic bonding involves the complete transfer of one or more electrons from a metal to a nonmetal, often giving both elements a noble-gas configuration.
In ionic compounds, electrostatic forces hold the ions together in crystal lattices.
Covalent bonds form through sharing of electrons between atoms.
Lewis structures represent the arrangement of bonding and unshared valence electrons in molecules and ions.
Two atoms can share one electron pair to form a single bond, two electron pairs to form a double bond, or three electron pairs to form a triple bond.
When there is more than one way to draw a Lewis structure, the molecule exhibits resonance. The actual bonding is a composite of all the resonance structures, and two or more Lewis structures are needed to show the arrangement of electrons in a covalent molecule or ion.
Carbon atoms bond together readily to form chains of atoms that are surrounded by hydrogen atoms to form the hydrocarbons.
Hydrocarbons can be classified as aliphatic or aromatic.
Aliphatic hydrocarbons can be further classified as alkanes, alkenes, or alkynes, depending on the types of bonds between carbon atoms.
Aromatic hydrocarbons contain rings of six carbon atoms with delocalized bonds.
Many other carbon compounds are formed from hydrocarbons by the inclusion of functional groups.
VSEPR theory, which assumes that electron domains stay as far apart as possible, predicts the shapes of molecules.
The total number of bonded atoms and of unshared electron pairs on the central atom determines the parent structure.
The molecular shape is determined by the positions of atoms around the central atom, which is affected by the presence of unshared electron pairs.
A nonpolar molecule has either nonpolar bonds or a symmetrical structure.
A molecule is polar if the bonds are polar and the molecule is not symmetrical.