We are skipping Chapter 6 on Thermochemistry, since you will explore this topic extensively next semester. For a preview, all we will cover is the meaning of delta H, enthalpy.
Reaction- the formation and breaking of chemical bonds
Bonds- the transfer or sharing of electrons
Electrons (abbreviated e-) what are they? how were they discovered?
William Crookes used an evacuated tube hooked up to a source of electricity in 1879 and saw a fluorescent beam (see Fig. 7.1)
JJ Thompson in 1897 determined the particles causing this fluorescence were negatively charged based on their behavior in the presence of a magnetic field
Thompson found the fluorescence was independent of the identity of the residual gas in the tube....the particles must therefore be a property of all matter
Thompson determined the mass to charge ratio (me/e = -5.686x10-12kg/C) (see Fig 7.2) and decided (based on previous knowlege with H+) he had one of three situations:if the charge were similar to H+, then mathematically the mass of the particle must be much less than that of H+
or if the mass were similar to H+, then mathematically the charge must be much greater than that on H+
or something in between the two extremes
he suspected the first case to be true, but could not prove it
1909 Millikan's oil drop experiment (see Fig. 7.3) showed that the rate of fall (velocity) of a charged oil drop could be varied in the presence of an electric field such that the charge could only in multiples of a fundamental unit (e=-1.602x10-19)
Thompson knew the e- had to be neutralized by + charges, but was unsure of the arrangement and devised the "raisin pudding" model (see Fig. 7.4)
Ernest Rutherford (a student of Thompson's) studied radioactivity, a phenomenon of unstable heavy atoms giving off radiation during disintegrationRutherford bombarded metal foils with alpha particles (He2+)...most went through, but some were scattered at odd angles...not explained by the Thompson model of the atom...can be explained with a nuclear model (see Fig. 7.5)
Prior to 1914, atomic numbers were given out simply by arranging the elements in order of mass...now that the number of protons and electrons could be determined and the mass of them calculated, there was some additional mass to be explained. A neutral particle was postulated, the existence of which was proven in 1932 by Chadwick
A Crookes tube can also be constructed to detect positive ions drawn to the cathode (negative pole) (or deflected away from the anode (positive pole) (again refer to Fig. 7.1 and also 7.6).
These experiments ARE matter dependent
A mass analyzer (mass spectrometer) is based on this concept expressed in a great analogy on p. 266-7
Electromagnetic (EM) radiation
Composed of perpendicular electric and magnetic field waves (something that repeats as it progresses through space)
The region of the full spectrum which WE call light is only a small portion (see Fig. 7.10&11)
Wavelength (lambda) is the distance between equivalent wave points on adjacent peaks, expressed in meters
Frequency (nu) is how many waves pass a fixed point per unit time, expressed in hertz (Hz, s-1)
Frequency times wavelength = speed, c= 2.99792458 x 108 m/s (or just 3 x 10 8 m/s)
In 1900, Max Planck explained the behavior of both high and low frequency radiation by equating energy with frequency and a constant, E=h(nu), where h is Planck's constant = 6.626 x 10-34 (think how small this number is)
Planck's equation demonstrates that light comes only in discrete "packets" called quanta (singular = "quantum")
In 1905, Einstein used Planck's theories to explain how electrons can result from bombarding a sample with light (the photoelectric effect, see Fig. 7.15)....this plus the observation of line spectra when elements were heated (see Fig. 7.12, 7.13,& 7.14 in Organic Lab) lead to.....
The Bohr model of the atom
Combined classical physics and quantum theory
Different energy levels of electrons correspond to orbits of different distances from the nuclei of atoms
The lowest energy level (that nearest the nucleus) is level 1, the next is level 2, and so on
The electron energy levels En = -B/n2, where n is the energy level (an integer) and B is a constant related to Planck's and the mass and charge of an electron (negative energy for an attractive force)
Bohr then explained line spectra as being due to the energy difference between 2 levels 
(see Fig 7.17&18)
When an atom has all its electrons in their lowest possible energy levels, the atom is in its ground state
If energy has been supplied sufficient to promote an electron to a higher level, the atom is in an excited state
If light can act like particles of matter, can particles act like waves? According to DeBroglie's theories, 
, yes!
The Bohr model is mostly classical, but if the particles are treated as waves, quantum or wave mechanics are needed
In 1926 Erwin Schrodinger developed a mathematical equation to describe the hydrogen atom (a wave equation, the solution to which is called a wave function
According to Max Born, the square of the wave function (psi2) gives the probability of finding an electron in a particular volume of space in an atom
According to Werner Heisenberg, in fact, we cannot know the exact position and motion of a tiny particle like an electron simultaneously....think of it this way, the act of measuring its position changes its motion, and vice versa (see Fig. 7.22)
Quantum numbers are integral value parameters from the wave function of the hydrogen atom...a set of these three wave function quantum numbers (there is a fourth) is called an atomic orbital, a mathematical expression which allows us to visualize a 3-D region in an atom where there is a significant probability of finding an electron (see Table 7.1 and Fig. 7.24, 7.25)
Principal quantum number (n), 1,2,3,4,5,....
Orbital angular momentum quantum number (l), 0,1,2,3,...(n-1) (s,p,d,f)
Magnetic quantum number (ml), 0, +/-1, +/-2,...+/-l(see Fig. 7.26, 7.27)
Spin quantum number (ms), +1/2, -1/2(see Fig. 7.28, 7.29)