Download Review II

yes no Was this document useful for you?
   Thank you for your participation!

* Your assessment is very important for improving the work of artificial intelligence, which forms the content of this project

Document related concepts

Theoretical and experimental justification for the Schrödinger equation wikipedia, lookup

Matter wave wikipedia, lookup

Hydrogen atom wikipedia, lookup

Molecular orbital wikipedia, lookup

Molecular Hamiltonian wikipedia, lookup

Ferromagnetism wikipedia, lookup

X-ray fluorescence wikipedia, lookup

Wave–particle duality wikipedia, lookup

Ionization wikipedia, lookup

X-ray photoelectron spectroscopy wikipedia, lookup

Auger electron spectroscopy wikipedia, lookup

Tight binding wikipedia, lookup

Atomic orbital wikipedia, lookup

Bohr model wikipedia, lookup

Chemical bond wikipedia, lookup

Electron configuration wikipedia, lookup

Atomic theory wikipedia, lookup

REVIEW CHAPTERS 3, 4, 5 (part) and 6
Atoms and Elements
Chapter 3
A. Atoms make up all matter
B. ~91 different naturally occurring elements
Indivisble: The Atomic Theory
A. Democritus (~400 BC)
B. Dalton (early 1800's)
1. Each element is composed of tiny
indestructible particles called atoms
2. All atoms of a given element have the same
mass and other properties that distinguish them from
atoms of other elements
3. Atoms combine in simple, whole-number
ratios to form compounds
-Law of conservation of matter
3.2 and 3.3 The Atomic Structure and Nucleus
A. Positive and negative charges (Thompson)-plum pudding model
B. Mostly empty space (Rutherford)
C. Current nuclear theory of the atom
1. Most of the atom's mass and all of its positive charge is contained in the
2. Most of the atom is empty space with small fast moving electrons
3. Number of protons = number of electrons in neutral atoms
3.4 The Properties of Protons, Neutrons, and Electrons
-Atomic Number = Number of protons
- Mass Number = # of Protons + # of neutrons
A. Masses expressed in atomic mass units (amu)
B. Protons and neutrons about 2,000 more massive than electrons
C. Nature of electrical charge
1. Electrical charge is a fundamental property of protons and electrons
2. Species of opposite electric charge are attracted to each other
3. Species of the same electric charge repel each other
4. Species of opposite electric charge cancel to make a neutral atom
Mass #
At. #
Isotopes: When the Number of Neutrons Varies in an element
A. Isotopes have the same chemical properties, but different masses
B. Some isotopes are more prevalent than others
C. Chemical symbol with mass number indicates which isotope
D. Some elements have many isotopes, some very few
Elements: Defined by Their Number of Protons
A. Identity of atom comes from the number of protons
B. Atomic number (Z) = number of protons
C. Chemical symbol is a shorthand notation of chemical name
1. C = carbon = 6 protons
2. O = oxygen = 8 Ptrotons
3. Co = cobalt = 27 Protons
Atomic Mass
A. Value on periodic table represents average mass of all isotopes
B. No chlorine atom weighs exactly 35.453 amu
3.5 and 3.6 Looking for Patterns: The Periodic Law and the Periodic Table
A. Mendeleev (1834 - 1907)
B. Periodic law
C. Metals (usually loose electrons to form cations)
D. Nonmetals (usually gains electrons to form anions)
E. Metalloids, also known as semiconductors
F. Individual group names
1. Group 1 – alkali metals +1 ion only
2. Group 2 – alkali earth metals
+2 ion only
3. Group 7 – halogens
usually –1 ion
4. Group 8 – noble gases
3.7 Periodic Trends: Ionization Energy, Atomic Size, and Metallic Character
A. Atomic size
1. Decreases as you move to the right along a row (Period)
2. Increases as you go down a column (group)
B. Ionization energy
1. Increases as you move to the right along a row (Period)
2. Decreases as you go down a column (group)
Chapter 4-Modern Model of Atom
4.1 Light: Electromagnetic Radiation
A. Light is not matter, it is pure energy
B. A photon is a particle of light
1. Wavelength is distance between adjacent wave crests
2. Frequency is the number of cycles that pass a point per unit time
The Electromagnetic Spectrum
A. Gamma rays
1. High energy
2. Low wavelength
B. X-rays
C. Ultraviolet
D. Visible
1. Very narrow band
2. Does not affect biological molecules
E. Infrared
F. Microwaves
G. Radio waves
4.2 Classical vs Quantum physics for matter and light
4.3 The Bohr Model: Atoms with Orbits
A. Electrons exist in quantized Circular orbits
B. Electrons get excited to a higher energy level by absorbing energy
C. Electrons emit light when relaxing from higher levels
D. Absorbed and emitted light correspond to energy difference between states
E. Higher the principal shell# higher is the energy of the electron (away from nucleus)
4.5 and 4.8 The Quantum-Mechanical Model: Atoms with Orbitals
A. Orbitals instead of Circular orbits
B. Electrons do not move in predictable paths
C. Probability maps-different shapes- called as ORBITALS
Quantum-Mechanical Orbits
A. Quantum numbers
1. Principal shell
2. Subshell: s (2e), p(6e), d( 10e), f(14e)
B. Electronic configuration
1. Gives location of each electron in the atom
2. Li: 1s22s1
C. How do the electrons fill in the atom
1. Auf-Bau Principle, Hund’s rule
2. Electron spin-Pauli exclusion principle
Filling order 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s-MEMORIZE
4.4 Electron Configurations and the Periodic Table
A. Valence electrons (outermost shell electrons with highest principal shell#)
B. Core electrons
C. Writing electronic configurations
1. Noble gas core
2. Locations of valence electrons
3. Highest n value is equal to row value
4. d electron group is equal to row number – 1
4.6 Compound Formation
A. Valence electrons dictate chemical reactivity
B. Elements in the same group have similar reactivity as they have
C. Compound formation and Octet rule with Chapter 5
4.7 Periodic Trends: Ionization Energy, Atomic Size, and Metallic Character
A. Atomic size
1. Decreases as you move to the right along a row (period)
2. Increases as you go down a column (group)
Molecules and Compounds
Chapter 5 Except section 5.7
5.1 Sugar and Salt
A Molecular View of Elements and Compounds
A. Atomic elements exist in nature as single atoms
B. Molecular elements exist as diatomic molecules, H2, N2, O2, F2, Cl2, Br2, I2
5.2 Compounds held by covalent bonds
-Molecular compounds formed between two or more nonmetals
- Attractive forces maximum at bond position
- Sharing of electrons
A. Every sample of a compound has the same proportions of constituent elements
5.3 Representing Valence Electrons with Dots (LEWIS SYMBOLS)
A. Valence electrons in outermost shell electrons
B. Lewis Structure is simply chemical symbol + # valence electrons around it
C. Most atoms "want" eight electrons around them (octet rule)
D. Exception: duet rule for hydrogen and helium
E. A covalent bond is a shared pair of electrons to get an octet around each atom
5.3 and 5.4 Covalent Lewis Structures: Electrons Shared
A. Two bonded nonmetals share electrons such that both get an octet (duet)
B. Both species get "credit" for all electrons in bond
C. Two species can share two, four, or six electrons
1. Two electrons shared is a single bond
2. Four electrons shared is a double bond
3. Six electrons shared is a triple bond
Writing Lewis Structures for Covalent Compounds
A. Write the correct skeletal structure for the molecule
B. Calculate the total number of valence electrons
C. Distribute electrons among atoms giving each an octet (duet)
D. If any atoms the lack an octet, form double or triple bonds as necessary
E. Exception to the octet rule
1. Duet rule for hydrogen
2. Odd number of electrons
3. Boron usually has only six electrons about it
4. Some elements in the third row and beyond may have more than
eight electrons around them.
A. Ionic compounds formed between a metal and one or more
B. Groups of atoms are set off by parentheses, e.g. Mg(NO3)2
Writing Formulas for Ionic Compounds
A. Ionic compounds always contain positive and negative ions
B. Formula units form neutral species
5.6 Electronegativity and Polarity:
A. Electronegativity
B. Bond dipoles
C. Pure covalent bonds
D. Polar and nonpolar bonds
E. Ionic bonds
6.2 Predicting the Shapes of Molecules (around central atom)
A. VSEPR Theory
1. Electron groups repel each other
2. Electron geometry
a. Linear
b. Trigonal planar
c. Tetrahedral
B. Predicting structures
1. Draw the Lewis structure
2. Count the number of electrons groups (Steric number)
3. Determine the number of bonding groups and lone pairs
4. Determine molecular geometry
6.3 Polarity of molecules
-Add up dipoles, if everything cancels, NON polar (m = 0)
- Dipoles do not cancel out – Polar molecule ((m ≠ 0)
-Lone pair on central atom usually causes molecules to be polar
- Unsymmetrical molecule- POLAR
Sample Questions for Second Test
Which of these elements have the highest electronegativity?
(a) Al
(b) S
(c) O
(d) B
Give the electronic configuration and outer shell configuration (i.e. 4s24p1) of the
following elements and state the number of valence electrons (i.e. 3). (DO NOT
use the shorthand “noble gas” notation)
Full Configuration
Valence config.
Number of valence eNa
Draw the Lewis structure, give the geometry of the molecule, and state
whether the molecule is polar (has a dipole) or nonpolar for the following:
Lewis Structure
1. CO2