Chemistry Review (Chapter 7-9)

Cathode rays are formed when electricity passes through gases at very low pressures.

Thomson's extensive studies of cathode rays proved taht they are negatively charged fundamental particles (which are now called electrons) found in all matter.

Thomson also established the mass-to-charge ratio of electrons by observing their deflections in electric and magnetic fields.

The charge on an electron was established by Millikan in a study of behavior of electrically charged oil droplets in an electric field.

An electron bears one fundamental unit of negative electric charge.

Thomson's raisin pudding model viewed an atom as electrons ("raisins") embedded in a positively charged "pudding."

Rutherford's interpretation of the scattering of alpha particles by metal foils suggested that the atom had a tiny, massive, positively charged nucleus that is surrounded by electrons.

Later experiments by Rutherford and others confirmed his atomic model and showed that the nucleus consists of protons and neutrons, which account for practically all the mass of an atom.

Weighted atomic masses of the elements and relative abundances of the isotopes of an element can be established by mass spectrometry, a technique that seperates charged particles according to their mass-to-charge ratios.

Electromagnetic radiation is an energry transmission in the form of oscillating electric and magnetic fields.

Oscillations produces electromagnetic waves that are characterized by their frequencies, wavelengths and velocity.

The complete span of possibilities for frequency and wavelength is described as the electromagnetic spectrum, of which visible light is only a very small part.

Light consisting of essentially all wavelength components in a region of the electromagnetic spectrum is a continuous spectrum (for example, the rainbow of colors obtained when a beam of sunlight is dispersed by a glass prism or raindrops in the sky).

Energetically excited gaseous atoms typically emit an emission spectrum containing only a discrete set of wavelength components in the form of a line spectrum.

Planck explained the spectral characteristics of black-body radiation by proposing that the energy of electromagnetic radiation is limited to intergral multiples of fundamental quantify that is called a quantum and has the value E=hv where h is defined as Planck's constant.

The photoelectric effect is the emission of electrons by illuminated surfaces. This is explained by thinking of quanta of energy as concentrated into particles of light called photons.

Bohr's theory requires the electron in a hydrogen atom to be in one of a discrete set of energy levels. The fall of an electron from an excited state to a ground state releases a discrete amount of energy as a photon of light with a characteristic frequency. Bohr's theory accounts for the observed atomic spectrum of hydrogen and is also applicable to the spectra of other one-electron species, such as He⁺ and Li²⁺.

A more satisfactory quantum mechanical treatment of atmoic structure is based on two additional fundamental concepts. These concepts are the wavelike properties of small particles and the Heisenberg uncertainty princeiple. These concepts transform the precisely described Bohr atom into an atomic model that replaces certainty with probability. Specifically, De Broglie proposed that a particle of mass moving at speed will have a wavelength. Thus, the electron in a hydrogen atom can be viewed as a matter wave enveloping the nucleus.

The representation of the wavelike properties of the electron through a wave equation is called quantum mechanics, and solutions of the wave equation are wave functions.

Wave functions require the assignment of three quantum numbers: the principal quantum number, n, the orbital angular momentum quantum number, l, and the magnetic quantum number m_l.

Atomic orbitals are wave functions with acceptable values of the three quantum numbers.

The wave mechanical treatment of the hydrogen atom can be extended to multielectron atoms, but with two essential differences: (1) Energy levels are lower in multielectron atoms than in the hydrogen atom, and (2) the energy levels in multielectron atoms are split; that is, the different subshells of a principal shell have different energies. The order of increasing subshell energy in a principal shell is s<p<d<f. All the orbitals within a subshell, however, have the same energy (they are degenerated orbitals).

Electron configuration is the distribution of electrons in orbitals among the subshells and principal shells in an atom. Three types of notation of electron configuration are the spdf notation, the noble-gas-core abbreviated notation, and the orbital diagram.

To write a probable electron configuration, we note that (a) electrons tend to occupy the lowest-energy orbitals available, (b) no two electrons can have all four quantum numbers alike (Pauli exclusion principle), and (c) where possible, electrons occupy orbitals singly and with parallel spins, rather than in pairs (Hund's rule).

The aufbau principle describes a process of hypothetically building up one atom from the atom having the preceding atomic number. With this principle and the other ideas cited in the chapter, it is possile to predict probable electron configurations for many of the elements. In the aufbau process for the main-group elements, electrons ar eadded to either the s or the p subshell of the principal shell of highest principal quantum number (n). In transistion elements, electrons go into the d subshell that is one shell in fromt he outermost occupied principal shell.

Elements with similar electron configurations fall in the same group of teh periodic table. For main-group elements, the group number correspons to the number of electrons in the principal shell of highest quantum number. The period number is the same as the principal quantum number of the valence shell. the division of the periodic talbe into s-, p-, d-, and f-blocks assists in the assignment of probable electron configureations. With in the f-block, the lanthanide series consists of elements in which the 4f subshell is the one being filled; the actinide series consists of elements in which the 5f subshell is being filled.

Valence electrons occupy the outermost occupied principal shell and have the highest principal quantum number of the electrons within the atom or ion.

Core electrons occupy the inner shells, those with principal quantum numbers less than that of the valence shell.

An atom with all electrons paired is diamagnetic, and an atom with one or more unpaired electrons is paramagnetic.

The periodic law describes the regular reccurrence of certain physical and chemical properties of atoms when they are arranged by increasing atomic number. Clear trends are seen in plots of atomic radius versus atomic number.

The specific definition of radius dimensions depends on the bonding environment in which elements are found; in molecules, we speak of the covalent radius of the atoms, and in metals, we speak of the metallic radius of the atoms. For ionic species, the size is expressed in terms of the ionic radius. Various species can be isoelectronic (contain the same number of electrons) but still have different atomic or ionic radii because the effective nuclear charge is different in each.

The effective nuclear charge (Z_eff) felt by an outer electron in an atom is the actual nuclear charge less the screening effect of all other electrons in the atom.

Ionization energy is the energy required to remove an electron from a ground-state atom (or ion) in the gaseous state.

Electron affinity is the energy change associated with the addition of an electron to a gaseous atom or ion.

The regions of the periodic table is designated to metals, nonmetals, metalloids, and the noble gases. This is related to the values of atomic properties. In general, matallic properties are associated with the ease with which atoms lose electrons; and nonmetallic properties, with the ease with which they gain electrons. The most metallic elements are located towards the left and bottom of the periodic table, and the most nonmetallic elements are located towards the right and top of the table.

The colors produced from metal flame test, oxidation and reduction properties, and the acid-base behavior of oxides all exhibit trends consistent with the periodic law. For example, oxides of nonmetals are generally acidic oxides; that is, they produce an acid upon reaction with water. Basic oxides, oxides that react with water to form bases, are generally oxides of metals. Oxides of elements having both metallic and nonmetallic character are amphoteric; they can react with either an acid or a base.

Chemical bonds form when the attractive forces between engatively charged electrons and positively charged nuclei exceed the repulsions between the nuclei to the maxiimum extent possible. The nature of the chemical bonds in a material determines many of its properties.

A Lewis symbol represents the valence electrons of an element, shown through dots distributed about the chemical symbol of the element. Lewis symbols of main-group elements are related to their locations in the periodic table. In forming compounds, main-group elements generally tend to follow an octet rule, which states that bonded atoms tend to acquire electron configurations having eight electrons in the valence shell.

An ionic bond forms through the transfer of electrons from metal to nonmetal atoms and the subsequent electrostatic attraction between the resulting ions.

An ionic crystal consists of an extended regular arrangement of ions.

The ionic bonds between nonmetals and s-block, p-block, and a few d-block metals can be represented through the Lewis symbols of respective anions and cations. this representation indicates the nature of electron transfer and the basis for electrostatic attration.

The net energy decrease in the formation of an ionic crystal from its gaseous ions is the lattice energy. A Born-Haber cycle relates the lattice energy and enthalpy of formation of an ionic compound, together with other atomic and molecular properties.

A covalent bond forms when two atoms share electrons.

A Lewis structure of a molecule show sthe arrangement of atoms in the molecule and covalent bonds between them. The lewis structure depicts electron pairs in the molecule as either bonding pairs or lone pairs. Usually, each atom in a Lewis structure acquires a noble-gas electron configuration, which for most atoms is a valence-shell octet of electrons.

A convalent bond involving one electron pair is called a single bond. A double bond involves two shared electron pairs, and a triple bond involves three.

Electronegativity (EN) is a measure of the ability of an atom to attract the electrons of a chemical bond to itself and is related to the position of an element in the periodic table. In a convalent bond between atoms of different EN values, electrons are displaced towards the atom with the higher EN. Chemical bonds vary from nonpolar covalent (zero or very small ∆EN) to polar covalent to ionic (large ∆EN).

Writing plausible Lewis strutures for molecules or polyatomic ions involves two general tasks: (1) writing the skeletal structure, which consists of one or more central atoms bonded to a number of terminal atoms, and (2) distributing the valence electrons among the atoms so as to (usually) follow the octet rule Resonance describes a phenomenon in which two or more Lewis structures have the same skeletal structure but different distributions of electrons among the bonded atoms. The best description of the actual structure, the resonance hybrid, is a combination of plausible resonance structures. In a resonance hybrid, some of the bonding electrons are delocalized over serval atoms. The resonance hybrid should conform to experimental data, such as bond lengths and/or bond energies, if these data are available.

Exceptions to the octet rule are found in odd-electron molecules and in molecular fragments calle dfree radicals. A few structures appear to have too few electrons to complete the octets of all atoms in the molecule or ion. Some structures appear to have too many. In the latter case, a central atom may have an expanded valence shell involving five, six, or even more elctron pairs.

Bond length, which in some cases can be related to atomic radii, is the internuclear distance between two bonded atoms. Bond length is greatly affected by bond order, that is, whether the bond is single, double, or triple.

Bond-dissociation energy is the energy required to break one mole of bonds in a gaseous species; its value depends on (1) the atoms in the bond, (2) the bond order, and (3) the particular molecule in which the bond is found. tabulated values are often averaged over a number of molecules containing the same bond. Bond-dissociation energies and average bond energies can be used to estimate enthalpy changes of reactions via the relationship ∆H = ∆H_{bonds broken} + ∆H_{bonds formed}.

Unsaturated hydrocarbons have one or more multiple bonds between carbon atoms. Alkenes have double bonds, and alkynes have triple bonds. A characteristic reaction of alkenes and alkynes is their ability to add atoms across the multiple bonds, thereby converting the multiple bonds to single bonds.

Some molecules containing multiple bonds undergo polymerization, a reaction in which small molecules (monomers) join together in large numbers to produce a gigantic molecule (polymer). Polymers typically are made up of long carbon chains, have high molecular massess, and often are represented through the structure of only the repeating unit.

Atomic Radius increases as the group decreases and the period increases in the periodic table.

Ionic Radius increases as the group decreases and the period increases in the periodic table.

Electron Affinity increases as the group increases and the period decreases in the periodic table.

Ionization Energy increases as the group increases and the period decreases in the periodic table.

Electronegativity increases as the group increases and the period decreases in the periodic table.