Chemistry Review
The Atom
• Nucleus:
–
–
: positive
: neutral
• Empty Space around Nucleus:
–
: negative
The Atom
The Atom
The Atom
• ATOMIC NUMBER (Z):
– Protons, Electrons
– Ex: Z = 10
Protons:
Electrons:
The Atom
• MASS NUMBER (A):
– Protons PLUS Neutrons
– Total number of particles in nucleus
The Atom
Also can be written as Carbon-14 or C-14.
Symbol of element – mass number
The Atom
• TO FIND:
– # of Protons: Atomic number
– # of Electrons: Atomic number
• If POSITIVE—subtract charge # from atomic #
• If NEGATIVE—add charge # to atomic #
– # of Neutrons: Mass number – atomic number
The Atom: Practice
• Fluorine-19
– # protons:
– # electrons:
– # neutrons:
• Uranium-238
– # protons:
– # electrons:
– # neutrons:
The Atom: Practice
•
4 He
2
– # protons:
– # electrons:
– # neutrons:
•
231 Th
91
– # protons:
– # electrons:
– # neutrons:
The Atom
• ISOTOPES: Atoms of same element with
different numbers of NEUTRONS.
• Different number of protons = different
elements.
The Atom
• Average atomic mass:
– Average of isotope weights
• Example: Fluorine
– 25 % F-19, 75% F-21
– (0.25 x 19) + (0.75 x 21)
The Atom: Practice
How many protons and electrons are in a 6429Cu2+
ion?
A 27 protons, 29 electrons
B 27 protons, 31 electrons
C 29 protons, 27 electrons
D 29 protons, 31 electrons
The Atom: Practice
How does an S2- ion differ from an electrically
neutral sulfur atom?
A.
B.
C.
D.
Mass number
Atomic number
Nuclear charge
Number of electrons
The Atom: Practice
• What are the differences between these isotopes
of hydrogen shown below?
1 H, 2 H, 3 H
1
1
1
A. The number of electrons and the atomic
number
B. The number of protons and the atomic number
C. The number of neutrons and the mass number
D. The number of protons and electrons
The Atom: Practice
• Which elements have the same number of
neutrons?
A. 105B and 126C
B. 5525Mn and 5626Fe
C. 10847Ag and 11248Cd
D. 19779Au and 20180Hg
Radioactivity
• Too many neutrons, not enough protons
• Atom wants to be STABLE.
• Releases particles/energy (radiation)
Radioactivity
Radioactivity
Radioactivity
Radioactivity
• Subtract the numbers on the particle symbols:
– 23592 U  42He + ?
– 235 – 4 = 231
– 92 - 2 = 90
? = 23190 Th
Radioactivity
• Half life: Time for half of original material to
decay to new stuff
•
•
•
•
•
Example : Half-life is 5000 years
Start: 100 grams
After 5000 years: 50 grams
After 10000 years: 25 grams
Etc…
Radioactivity
• Nuclear Reactions: Fission and Fusion
Radioactivity: Practice
When 4219K undergoes radioactive decay, the result is two
products, one of which is calcium-42. What is the other
product?
A. 42He
B. 24He
C. 11e
D. 01−e
Radioactivity: Practice
In which group are the particles arranged in
order of decreasing mass?
A. alpha, beta, neutron
B. alpha, neutron, beta
C. neutron, beta, alpha
D. neutron, alpha, beta
Radioactivity: Practice
The half-life of phosphorus-32 is 14.3 days. How
much of a sample of phosphorus-32 will remain
after 57.2 days?
A.
B.
C.
D.
1/32
1/16
1/8
1/4
Electromagnetic Spectrum
Electromagnetic Spectrum
• Radio wave: longest wave, lowest frequency
• Gamma ray: shortest wave, highest frequency
• Wavelength goes up, frequency goes down
• Wavelength goes down, frequency goes up
• Frequency goes up, energy goes up
The Electron
• EM energy: comes in small units known as
quanta
– Photon = quantum
• Atoms get energy  Electrons shoot up to
higher levels
• Atoms lose energy  Electrons fall back to
original level; energy released as photons
The Electron
The Electron
• Bohr’s Model:
Electrons in
discrete levels
aroud nucleus
• Do not move
from level
unless energy
gained or lost
The Electron
• Bohr model only useful for hydrogen
• Quantum model more accurate:
– Electrons act as both waves and particles
– Exist in a cloud around nucleus where they are
LIKELY to be
The Electron
The Electron
• First: Draw out levels, starting at bottom and
going up.
– 1s  2s  2p  3s  3p  4s  3d  4p  5s 
4d  5p  6s  4f  5d  6p  7s
• Remember:
–
–
–
–
S has 1 blank
P has 3 blanks
D has 5 blanks
F has 7 blanks
The Electron
The Electron
• Rules for Writing Configurations
– 1. Start at bottom. Work your way up.
– 2. Fill every blank in a row with one arrow before you
put a second arrow in each blank.
– 3. Must have one up arrow and one down arrow. Not
two up, not two down.
• Gives lowest energy configuration.
The Electron
• After drawing out configuration, use arrows to
write out configuration.
– Exp: 1s2 2s2 3s2
• To identify element, either count arrows or
count the “exponent” numbers.
– In above example, there is 2 + 2 + 2 = 6 electrons.
Element is carbon (atomic # = 6)
The Electron
• Draw and write out electron configuration for
silicon.
The Electron
The Electron
• Draw and write out electron configuration for
magnesium.
The Electron
The Electron
• Blocks
nickel
D
lead
The Electron: Practice
EL
8.
C
Which orbital notation represents an s-block element in the third period?
A
1s
2s
1s
2s
2p
3s
1s 2s
2p
3s
3p
2s
2p
3s
3p
R
B
C
D
1s
Page 2
4s
3d
Go t o
The Electron: Practice
• Which electron transmission in the hydrogen
atom will result in the emission of red light?
A.n = 2 to n =3
B.n = 2 to n = 4
C.n = 3 to n = 2
D.n = 4 to n = 2
The Electron: Practice
In which block does an element with the
electron configuration [Xe] 6s2 4f14 5d10 6p1
belong?
A.s block
B.p block
C.d block
D.f block
The Electron: Practice
Nor t h Car ol ina Test of Chemist r y. For m A REL EASED Fall 2009
ion?
r eact
react
Which orbital notation shows the
lowest energy arrangement of valence
electrons for 1s22s22p3 ?
A
B
C
D
2s
2s
2s
2s
2p
D
51.
2p
2p
The Periodic Table
The Periodic Table
• Rows: Periods
• Columns: Groups
– Groups 1, 2, 13-18: Representative elements
– Groups 3-12: Transition metals
The Periodic Table
The Periodic Table
• METALS:
– High boiling, high melting points
– Conduct electricity and heat
– Shiny, ductile
– As you go DOWN a group, metallic nature of
elements INCREASES
The Periodic Table
• NONMETALS:
– Low melting, low boiling points
– Poor conductors of electricity and heat
– Brittle
The Periodic Table
• Group 1: Alkali metals
• Group 2: Alkaline earth metals
• Group 7: Halogens
• Group 8: Noble Gases
The Periodic Table
• ATOMIC RADIUS: How big an atom is.
– Down a column  Radius INCREASES
– Left to right across a row  Radius DECREASES
The Periodic Table
• ELECTRONEGATIVITY: A measure of how well
an atom can attract electrons to itself
– Up a column: INCREASES
– Left to right: INCREASES
The Periodic Table
The Periodic Table
• IONIZATION ENERGY: Energy needed to
remove one electron from an atom.
– Down a group: DECREASES
– Left to right: INCREASES
Bonding
• VALENCE ELECTRONS: Electrons in the
outermost level
– # of Valence Electrons = Group number for
Representative Elements
• OCTET RULE: Atoms are stable when they
have 8 valence electrons
Bonding
• ION: a charged atom
• CATION: a positive ion
– Losing electrons and size  Cation SMALLER than
neutral atom
• ANION: a negative ion
– Gaining electrons and size  Anion LARGER than
neutral atom
Bonding
• Metals make cations
• Example: Sodium  Group 1 1 valence
electron
– Easier to lose 1 electron to get 8
– Lose 1 electron  +1 charge
Bonding
• Nonmetals make anions
• Example: Fluorine  Group 7
– Has 7 valence electrons
– Easier to gain 1 to get 8
– Gain 1 electron  -1 charge
BONDING
• IONIC BONDING: a cation and an anion bond
together
• The cation gives up its electron to the anion
• Metal + Nonmetal
Bonding
Bonding
• REMEMBER:
– Take the charges and swap them to find the
bottom number
– Magnesium  Mg2+
– Chlorine  Cl-
– Mg2+ Cl-  MgCl2
Bonding
• Naming Ionic Compounds:
– First part (Cation): Just say the name
– Second part (Anion): Change the ending to “-ide”
– Example: MgCl2
• Magnesium Chloride
Bonding
• Polyatomic Ions
– More than 1 atom  Treat as one ion
– Page 7 in Reference Table
Bonding
• COVALENT BONDING: Atoms will SHARE
electrons
• Two nonmetals
Bonding
Bonding
• Naming Covalent Bonds:
– Use the prefix system (Mono, di, tri, tetra, penta,
etc..)
– Name first element as is (with prefix if necessary)
– Name second element with ending “-ide” (use
prefix if necessary)
Bonding
• Example: CO2
– 1 Carbon  Just carbon
– 2 Oxygen  dioxide
• Carbon dioxide
• Example: B3Fl5
– 3 Boron  Triboron
– 5 Fluorine  Pentafluoride
• Triboron pentafluoride
Bonding
• Ionic Compounds:
– Very strong
– Conduct electricity
• Covalent Compounds:
– Not as strong
– Poor conductors