Download Name January 5, 2017 Period Bio-Chem Unit 2 Review (Chapters

Name ___________________________________ January 5, 2017 Period _______________
Bio-Chem Unit 2 Review (Chapters 3&4)
I can……..
□ explain the of the Atomic Theory, and how it developed from the time of the Greeks to the early
1900’s.
□ explain the Law of Conservation of Mass/Energy, the Law of Definite Proportions, and Law of Multiple
Proportions.
□ describe contributions by Rutherford, Bohr, and Max Plank
□ find the number of protons, neutrons and electrons using the periodic table
□ understand that the average mass number is the average of the atomic masses of the naturally occurring
isotopes.
□ explain the trend in the periodic table, as one goes down each Group, (column) what is changing?
As you go across a Period, (row) what trends can be seen, what is changing?
□ calculate wavelength or frequency using the speed of light equation C = λv or C = λf
□ calculate the energy of a photon of light using Planck’s constant. E = hv or E = hf
□ draw and explain Lewis dot structures and Bohr models of an atom.
□ explain emission spectrum and how they are useful in determining the probability of where an electron
might be in an atom. (POGIL: Electrons Energy, and Light)
□ explain the different quantum numbers and what they represent
□ write an electron configuration for a 5 elements on the periodic table.
□ Compare and contrast the electron configurations in the first, second, third and fourth periods of the
periodic table.
I. Atomic theory:
A. 400 BC
B. 400 BC
Aristotle: Matter is made of Air, Fire, Earth, & Water.
Democritus: matter is not continuous but is composed of tiny, discrete, indivisible particles
called “atomos,” (atoms).
C. 1808 John Dalton was first to publish Atomic Theory
a. All matter composed of indivisible particles called atoms, which retain their identity during
chemical reactions
b. All atoms of the same element have identical properties, which differ from other elements.
c. Atoms cannot be created nor destroyed or transformed into other atoms of another element.
(except by nuclear reactions)
d. Compounds are formed when atoms of different elements combine with each other, in small
whole-number ratios.
e. The relative kinds and numbers of atoms are constant in a given compound.
D. 1897
JJ. Thompson: discovered the charge of the electron with his Cathode Ray Tube, and theorized
that the atom was a “plum pudding” of electrons and positive charges, (not separated by a
nucleus).
E. 1897-98
Rutherford: Used the Gold Foil experiment to prove the Plum Pudding Model, only to discover
and disprove that the atom has a nucleus filled with protons, and other significant mass. The atom
is mainly empty space, where the mass of the atom lies within the nucleus.
F. 1913
Niels Bohr: Electrons move about the nucleus in “orbits” in Quantum Energy levels. Tested this
in the hydrogen atom. The maximum number of electrons per energy level is, where n = the
energy level number
G. 1922
Max Plank: realized that light and other electromagnetic waves were emitted in discrete packets
of energy that he called "quanta" - "quantum" in the singlular - which could only take on certain
discrete values (multiples of a certain constant, which now bears the name the “Planck constant”).
This is generally regarded as the first essential stepping stone in the development of quantum
theory, which has revolutionized the way we see and understand the sub-atomic world
II. Law of Definite Proportions: (applies to compounds)
1. a compound is a pure substance consisting of two or more different elements in a fixed ratio.
2. Law of Definite Proportions states: different samples of any compound contain the same
element in the same proportion by mass.
Example: water is always found to have definite proportion 88.9% oxygen and 11.1%
hydrogen by mass. Why?
WHY? Because water is composed of particles in which 1 atom of O is attached to 2
atoms of hydrogen. Oxygen weighs 16 times as much as hydrogen.
III. Law of Multiple Proportions
Law of Multiple Proportions states when two of more elements form more than one compound with each
other, the masses of each element are in fixed small whole number ratios.
EXAMPLE: when H and O combine they can form to form H2O, (water) and H2O2 (dihydrogen
dioxide) in the same sample.
Amount of Oxygen per gram per gram of Hydrogen in H2O
Amount of Oxygen per gram per gram of Hydrogen in H2O2
= whole number ratios
= 8.0 grams
= 1 This is true that compounds combine in fixed whole number ratios
16.0 grams
2
IV. Structure of the Atom
New atomic theory: atoms are not the smallest particles. They are composed 3 fundamental parts:
a. electrons (e-): discovered by JJ Thomson 1897 (cathode ray tube).
b. protons (p+): discovered by 1897-98 Goldstein/Rutherford (cathode ray tube).
c. neutrons (n0): discovered by Chadwick 1932
V. Atomic mass or amu: was determined by taking the average atomic mass of all the naturally occurring isotopes of the
element.
VI. Atomic # = number of protons
Atomic mass number = number of protons + neutrons
To find number of neutrons subtract atomic mass – atomic number
In an electrically neutral atom: protons = number of electrons
VII. Coulombic law:
1. like charges repel, opposites attract
2. larger the distance between protons and electrons less of an attractive charge.
3. the more protons in the nucleus the greater the attractive charge
VIII. See Bell Work from last week and POGIL: Electron Energy and LIGHT
Calculate wavelength or frequency using the speed of light equation C = λv or C = λf
Calculate the energy of a photon of light using Planck’s constant. E = hv or E = hf
Calculate wavelength or frequency using the speed of light equation C = λv or C = λf
Calculate the energy of a photon of light using Planck’s constant. E = hv or E = hf
IX. Have some sense of what quantum numbers, and that there are 4 quantum numbers
n- Principal quantum number: represents the energy level the electron is in, linked to the periods of the periodic
table.
l- Secondary quantum number : shape of the orbital- s, p, d, f
ml- Magnetic quantum number: relates to quantum number l, how many (sub)orbital there are in each quantum
letter l for instance s = 1, p = 3 d = 5 and f = 7
+1/2, -1/2 Spin quantum number has only two values and has to do with Pauli Exclusion Principle = no two
electrons I the same orbital and indicates the spin of an electron.
X. Difference between emission/absorption spectrums and what they are depicting of electrons. These spectrums are
unique to each element like a fingerprint.
Name ______________________________________ Period ________________
Please write down the answers on a separate piece of paper.
Practice Problems: Due Friday. Be ready to ask questions before we take the test.
1. The total number of orbitals that can exist in the third main energy level is?
2. Dalton’s atomic theory helped to explain what law?
3. What happens when an electron goes from its ground state to a higher energy level?
4. Most of the volume of an atom is where?
5. A single orbital in the 3d level can hold how many electrons?
6. speed of light = wavelength x frequency c = λf
speed of light = c = 3.00 x 108 m/s
1nm = 10-9 m
E = hf
109 nm = 1m
h = Planck’s constant = 6.63 x 10-34 J·s
102 cm = 100cm = 1m
Solve the following problems using the equations above. First write the equation you are using.
Remember any conversions you need to complete before you compute.
a. What is the frequency in hertz of red light having a wavelength of 710 nm?
b. Ozone protects the earth's inhabitants from the harmful effects of ultraviolet light arriving from the
sun. This shielding is a maximum for UV light having a wavelength of 295 nm. What is the frequency in
hertz of this particular wavelength of UV light?
c. Radar signals are also part of the electromagnetic spectrum in the microwave region. A typical radar
signal has a wavelength of 3.19 cm. What is the frequency in hertz?
7. Principal quantum number 2 the total number electrons are?
8. True or False: atomic mass on the periodic table is the average atomic mass of all the natural occurring
isotopes of that element on earth.
9. The mass of a neutron is relatively the ___________________ as that of a(n) ___________________.
10. The majority of the mass of an atom is found in the ___________________________.
11. What are isotopes?
12. Define ion in your own words.
13. Find the number of neutrons for the following elements: Mg, Zr, P.
14. The number of orbitals for d are ?
15. Visible light, infrared, ultra-violet all have what in COMMON?
16. Ag-108 has how many neutrons?
17. The distance between two peaks of a wave is called what? (frequency, wavelength, speed, or energy)
18. Explain why an atom is electrically neutral and an ion is not.
19. What is the electron configuration for the following: Ni, N, Xe?
20. Name 3 types of electromagnetic radiation and put them in order from highest energy to lowest.
21. Draw a Bohr model of calcium.
22. Fill in the following chart:
Symbol
Atomic #
Protons
Neutrons
Mass number Average atomic mass
15.99
12
Na
69.723
12