CH. 4 Atomic Structure

Atoms

• *Static electricity demonstration (pg. 96 OB)*• Atom = the smallest particle of an element that

retains its identity in a chemical reaction• Rutherford discovered alpha, beta, and gamma

radiation and also positive charges and most mass is concentrated in the core of an atom

• The radii of most atoms fall b/w 5 x 10-11 m to 2 x 10-10 m but can be seen with a scanning tunneling microscope (STM)

Subatomic Particles

• 3 kinds include electrons, protons, neutrons• Electron – mass is so small it is considered irrelevant,

so atomic mass is the sum of protons and neutrons• Proton – positively charged subatomic particle• Neutron – no charge but w/ a mass nearly equal to

the proton

Rutherford Atomic Model

• Nucleus = tiny central core of an atom composed of protons and neutrons

• In Rutherford’s “Nuclear Atom” the protons and neutrons are located in the nucleus and the electrons are distributed around the nucleus and occupy almost all the volume of the atom

• Periodic trend – 1-20 equal p+ & n⁰, above 20 need more n⁰ to stay stable (strong nuclear force), 83 & above not enough n⁰ to stay stable

A neutron walks into a restaurant and orders a couple of cokes. As she is about to leave, she asks the waiter how much she owes. The waiter replies, "For you, No Charge!!!"

Atomic Number

• Elements are different because they contain different numbers of protons

• Arranged on the periodic table in numerical order by the proton number

• Atomic number = the number of protons in the nucleus of an atom of that element

• Atoms are electrically neutral so protons equal the number of electrons

Mass Number

• Mass number = the total number of protons and neutrons in an atom

• The number of neutrons in an atom is the difference b/w the mass number and atomic number

• Number of neutrons = mass number – atomic number N = M – P, P + N = M, P = M – N

Isotopes

• Isotopes = atoms of the same element (same number of protons) w/ diff. # of neutrons

• B/C isotopes of an element have diff. #’s of neutrons, they also have diff. mass numbers

• some are radioactive (treating cancers and aging objects)

Atomic Mass• The actual mass of each of an atom is too small to be practical so you

compare the relative masses by using a reference isotope as a standard which is carbon-12

• Atomic mass unit (amu) = one twelfth of the mass of a carbon-12 atom– Each proton or neutron in carbon is about 1 amu

• In nature most elements occur as a mixture of 2 or more isotopes• Atomic mass = weighted avg. mass of the atoms in a naturally occurring

sample of the element– to find this you need the # of stable isotopes (H has 3), the mass of each

isotope, and the natural percent abundance– Multiply the mass of each isotope by its natural abundance, then add the

products– EX: C-12 has a mass of 12.000 amu and abundance of 98.89%, C-13 is 13.003

amu and occurs 1.11%. What is the atomic mass of carbon?

Probs

• Pp 15-24• Sample prob 4.1-4.2• Conceptual prob 4.2• Pgs. 111-117• *Edible atom activity* and/or *The atomic

mass of candium*• *Research one isotope and write down its

practical application*

Ch. 3 Material BellringerPart I

• 1) Write 367,000,000 in scientific notation rounded to the nearest tenth

• 2) How many sig figs are in the following numbers: A)361.00560 B)123,450,000,000 C) 0.000200406

• 3) Perform the following calculation rounded to the correct sig figs: 43.4560 – 28.62

• 4) Perform the following calculation rounded to the correct sig figs: 68.3200 / 4.200

• 5) Which set of numbers is more precise and which is more accurate if the accepted value is 68.32– A) 66.32, 66.98, 64.32 B) 68.12, 68.48, 68.23 C)70.12, 69.14,

68.95

Ch. 5 Electrons in Atoms

Bohr Model

• Rutherford could not explain the chemical properties of elements so his student Niels Bohr improved on it by proposing that an electron is found only in specific circular paths, or orbits around the nucleus

• energy levels = fixed energies an electron can have• The higher energy levels are closer together so it takes

less of a quantum of energy for an electron to move from one level to the next the farther you move from the nucleus

• Quantum = amount of energy required to move an electron from one energy level to another

Quantum Mechanical Model

• Electrons move around the outside of the nucleus on energy levels so fast it just looks like a cloud. The electrons aren’t seen while looking at the cloud but where they are likely to be found it is denser and is expressed by a probability.– 1. energy of electrons is quantized – only specific amts. of energy– 2. exhibit wavelike behavior– 3. impossible to know the exact position as well as momentum of

an electron at any given instant• Atomic orbital = a region of space in which there is a high

probability of finding an electron

Atomic Orbitals• For each principal energy level, there may be several orbitals w/diff.

shapes & at diff. sublevels• Each energy sublevel (s,p,d,f) corresponds to an orbital of a diff. shape,

which describes where the electron is likely to be found• s orbitals are spherical, p orbitals are dumbbell-shaped (3 have diff.

orientations in space), d most have cloverleaf shapes, f more complex

…cont

• Each orbital (or spot) can hold two electrons, one spinning clockwise and the other counter creating a magnetic field, thus each principal energy level can hold a certain max # of e’s

Electron Configurations

• Electron configurations = the ways in which electrons are arranged in various orbitals around the nuclei of atoms

• 3 Rules (tells how to find the e config.):– aufbau principle – electrons occupy the orbitals of lowest

energy 1st – Pauli exclusion principle – an atomic orbital may describe at

most 2 e’s. To occupy the same orbital, 2 e’s must have opposite spins; that is, the e spins must be paired

– Hund’s rule – e’s occupy orbitals of the same energy in a way that makes the # of e’s w/ the same spin direction as large as possible

aufbau principle

HundRule

• 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6, 6s2, 4f14, 5d10, 6p6, 7s2, 5f14, 6d10

• There are exceptions to the aufbau principle like chromium and copper b/c half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations. The higher you get, the closer the energy levels are and the more apt the e’s will jump to a more excited state

• Cp 5.1, pp 8, 9 pg. 135

Ch. 3 Material BellringerPart II

• 1) Calculate the percent error if a student working on a lab obtained the mass of a metal sample to be 94.3 g. The teacher told him he was wrong and to measure again b/c it was actually 94.8 g.

• 2) Convert 23 km to dm• 3) How many gallons are in 23,000 cm3 (3.785 L =

1 gal)

• 4) How many inches are in 38 cm (2.54 cm = 1 in.)• 5) Convert 40 mph to ft/sec (1 mi = 5280 ft.)

What led to the development of the quantum mechanical model of the atom?

LIGHT

Light

• Electromagnetic radiation = form of energy consisting of waves made up of oscillating electric and magnetic fields at right angles to each other– Includes: radio waves, microwaves, infrared

waves, visible light, ultraviolet waves, X-rays, and gamma rays

Waves• All waves have amplitude, wavelength, frequency, and speed• Amplitude = height of the wave measured from origin to its crest• Wavelength = distance b/w successive crests of the wave (λ)• Frequency = tells how fast the wave oscillates up and down, or # of waves

cycles to pass a given point per unit of time (ν)– Cycles per second – hertz (Hz) or 1/s or s-1 a reciprocal second

• Speed of light = 3 x 108 m/s (c) for constant speed• Wavelength and frequency of light are inversely proportional

– c = λ ν speed of light = wavelength x frequency• visible spectrum = portion of the EM spectrum that can be seen w/ the

unaided eye– ROYGBIV

• Red has lowest frequency, longest wavelength• Violet has highest frequency, shortest wavelength

• Sp 5.1 and pp 14-15 pg. 140

Atomic Spectra

• When atoms absorb energy, e’s move into higher energy levels. They then lose energy by emitting light when they return to lower energy levels– The light emitted by atoms consists of a mixture of

only specific frequencies. Each specific frequency of visible light emitted corresponds to a particular color. When the light passes through a prism, the frequencies of light emitted by an element separate into discrete lines to give the atomic emission spectrum of the element (the emission spectrum of an element is like a fingerprint)

Quantum Theory• When an electron has its lowest possible energy, the atom is

in its ground state (n=1)• Absorbing energy excites electrons to higher levels (n=2, 3, 4,

5, 6, …)• A quantum of energy in the form of light (photon) is emitted

when electrons drop back to a lower energy level• There is a restriction on amt. of energy that can be emitted or

absorbed– This quantum of energy E is related to the frequency of light

emitted• E = hv where E is energy, h is planck’s constant (6.6262 x 10-34 J-s) and v is

frequency• The light emitted by an electron moving from a higher to lower energy

level has a frequency directly proportional to the energy change of the electron

Why does hamburger have lower energy than steak? Because it's in the ground state.

Bellringer Answers

• Part I– 1) 3.7 x 10^8– 2) 8, 5, 6 – 3) 14.84– 4) 16.27– 5) B

• Part II– 1) .53%– 2) 230,000 dm– 3) 6.08 gal– 4) 14.96 in– 5) 58.67 ft/sec