Thursday, October 11 •  Lecture 11 (Wednesday) –  The Atom (3.3) –  Atomic Number & Mass Number (3.4) –  Isotopes and Atomic Mass (3.5) –  Electron Arrangement in Atoms (3.6) •  Lecture 12 (Today) –  Electron Arrangement in Atoms (3.6) –  Orbital Diagrams and Electron ConfiguraKons (3.7) –  Trends in Periodic ProperKes (3.8) Fixing Clicker Issues •  When you click in during class, make sure that your clicker shows a green light •  Check the list of working and non-­‐working clickers •  If you are in the “not working” column, re-­‐register your clicker •  double-­‐check that your Device ID has been submiXng responses (boYom of file) CorrecKons The 12C atom = 12 amu (exact) I said that the 12C nucleus = 12 amu (exact), which is wrong (even though the masses are basically the same) Calculating Atomic Mass for C
•  To calculate the atomic mass, we use a weighted average
•  The contribution to atomic mass is based on the
abundance and mass of each isotope
correct! isotope mass (amu) abundance contribu2on to atomic mass 12C 12 (exact) × 0.989 = 11.868 amu 13C 13.003 × 0.011 = 0.1430 amu 14C 14.003 × ~0 = 0 amu 12.0110 amu 12.0 amu •  The abundance values add up to 1 •  To convert the abundance to a % value, mulKply by 100 Some Isotope Comparisons Element Isotope Hydrogen 1H 2H Carbon 12C 13C Sulfur 32S 33S 34S 36S Lithium 6Li 7Li Mass of Isotope (amu) Rela2ve Abundance Average Atomic Mass (amu) 1.00782 2.01410 99.9844% 0.0156% 1.0079 12 (exact) 13.00335 98.892% 1.108% 12.01115 31.972071 32.971458 33.967867 35.967080 95.06% 0.74% 4.18% 0.0136% 32.064 6.015123 7.016005 7.5% 92.5% 6.941 5
Electromagne2c radia2on (“light”) is a form of energy, characterized by wavelength (λ). A galaxy imaged in the visible spectrum. The same galaxy imaged in the radio spectrum at the VLA. Scanning-­‐electron microscopy image X-­‐ray image visible spectrum infrared (IR) spectrum Atomic Emission and Absorbance Atoms can emit EM energy only at specific, discrete wavelengths. They can absorb EM energy only at these same discrete wavelengths. The energy that an atom can absorb/emit is quan<zed. QuanKzed vs. ConKnuous These balls can only stop on one of the stairs. This ball can roll all the way down the ramp. Their distance from the ground is quan<zed. Its distance from the ground can change con<nuously. Emission Spectrum of Hydrogen The discrete line spectra of atoms were explained as the movement of an electron between quanKzed energy levels. Principle quantum number, n
an index of the electron energy level M
n=5
n=4
n=3
The emi=ed photon shows up as one of the lines in the spectrum... n=2
As n increases… energy of the level increases distance from the nucleus increases distance between the levels decreases E
…but it represents a change between two energy levels in the atom. n =1
n ranges from 1 to ∞, but pracDcally speaking n = 7 is the highest level the electron can occupy and sDll be considered part of the atom. Sublevels
Within each energy level, we have sublevels that
§  contain electrons with identical energy.
§  are identified by the letters s, p, d, and f.
The number of sublevels within a given energy level
is equal to the value of the principal quantum number, n.
Orbitals
Simple Model: •  the electrons float around outside the nucleus Orbitals: •  3D regions where electrons are likely to reside s p d Orbitals
Each electron sublevel consists of orbitals, which §  are regions where there is the highest probability of finding an electron (ojen 90%). §  have their own unique three-­‐dimensional shape. §  can hold up to 2 electrons each. s Orbitals
We know that s orbitals have a spherical shape, centered around the atom’s nucleus. §  The s orbitals get bigger as the principal quantum number, n, gets bigger. §  Each s orbital can hold up to 2 electrons that must have opposite spins p Orbitals
There are three p orbitals in each energy level,
starting with energy level 2. They
§  have a two-lobed shape, much like tying two balloons
together, and can hold 2 electrons.
§  are labeled x, y, and z.
§  increase in size as n increases (n = principle energy level).
2px 2py 2pz d Orbitals
There are five d orbitals in each energy level,
starting with energy level 3. They
§  have a four-lobed shape, much like tying 4 balloons together,
and can hold 2 electrons.
§  are labeled xy, yz, xz, x2-y2, and z2.
§  increase in size as n increases (n = principle energy level).
3dxy 3dyz 3dxy 3dx2-­‐y2 3dz2 Levels, Sublevels, and Orbitals
•  Each level contains a specific number of sublevels
•  Each sublevel consists of a specific number of orbitals.
§ 
§ 
§ 
§ 
An s sublevel contains one s orbital.
A p sublevel contains three p orbitals.
A d sublevel contains five d orbitals.
An f sublevel contains seven f orbitals.
Order of Filling
Energy levels are filled with electrons
§  in order of increasing energy.
§  beginning with quantum number n = 1.
§  within an energy level, the
orbitals are filled in the order
s, p, d, f
n = 7 n = 6 n = 5 n = 4 n = 3 n = 2 n = 1 Orbital Diagrams
An orbital diagram shows
§  orbitals as boxes in each sublevel.
§  electrons in orbitals as vertical arrows.
§  electrons in the same orbital with opposite spins (up
and down vertical arrows).
Example:
Orbital diagram for Li
1s2 2s1 filled half-­‐filled 2p empty Writing Orbital Diagrams
The orbital diagram for
carbon has 6 electrons:
§  2 electrons are used to fill
the 1s orbital.
§  2 more electrons are
used to fill the 2s orbital.
§  1 electron is used in two
of the 2p orbitals so they
are half-filled, leaving one
p orbital empty.
Order of Filling
Electrons in an atom
§  fill the lowest energy
level and orbitals first,
§  fill orbitals in a
particular sublevel with
one electron each until
all orbitals are half full,
and then
§  This minimizes e−-e−
repulsion
§  fill each orbital using
electrons with opposite
spins
# e− 1s 2s 2p 1 H 2 He 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne Electron Configuration
An electron configuration
§  lists the filled and partially filled energy levels in order of
increasing energy.
§  lists the sublevels filling with electrons in order of
increasing energy.
§  uses superscripts to show the number of electrons in
each sublevel.
§  for neon is as follows: number of electrons = 10
1s22s22p6
Electron
Configurations
Electron configura2ons show •  the number of electrons •  which orbitals are occupied NOTE: only the far right porKon of the electron configuraKon changes between elements # e− element electron configura2on 1 H 1s1 2 He 1s2 3 Li 1s22s1 4 Be 1s22s2 5 B 1s22s22p1 6 C 1s22s22p2 7 N 1s22s22p3 8 O 1s22s22p4 9 F 1s22s22p5 10 Ne 1s22s22p6 Abbreviated Configurations
In an abbreviated configuration,
§  the symbol of the noble gas is in brackets, representing
completed sublevels.
§  the remaining electrons are listed in order of their
sublevels.
Example: Chlorine has the following configuration:
1s22s22p63s23p5
[Ne]
The abbreviated configuration for chlorine is
[Ne]3s23p5.
Period 2 Configurations
Period 3 Configurations
Learning Check
Which of the following are correct electron
configurations for Sulfur?
A.  1s22s22p63s23p4 B.  [He] 3s23p4 C.  [Ne] 3s23p4 D.  3s23p4 E.  3p16 F.  A and B G.  A and C H.  C and D Electron Configurations and the
Periodic Table
The periodic table consists of sublevel blocks
arranged in order of increasing energy.
§  Groups 1A and 2A
= s block
§  Groups 3A to 8A
= p block
§  Transition Elements
(This sublevel is (n-1), 1 less
than the period number.) = d block
§  Lanthanides/Actinides
(This sublevel is (n-2), 2 less
than the period number.) = f block
Sublevel Blocks
S Writing Electron Configurations
Using the periodic table, write the electron configuration
for silicon.
Solution:
Period 1 1s block
1s2
Period 2 2s → 2p blocks
2s2 2p6
Period 3 3s → 3p blocks 3s23p2 (at Si)
Writing all the sublevel blocks in order gives the
following:
1s22s22p63s23p2
Writing Electron Configurations
Using the periodic table, write the electron configuration
for manganese.
Solution:
Period 1 1s block
1s2
Period 2 2s → 2p block 2s2 2p6
Period 3 3s → 3p block 3s2 3p6
Period 4 4s → 3d block 4s2 3d5 (at Mn)
Writing all the sublevel blocks in order gives the
following:
1s22s22p63s23p64s23d5
Valence Electrons
The valence electrons
§  determine the chemical properties of the elements.
§  are the electrons in the outermost, highest energy level.
§  are related to the group number of the element.
Example: Phosphorus has 5 valence electrons.
5 valence
electrons
P Group 5A(15) 1s22s22p63s23p3
Groups and Valence Electrons
All the elements in a group have the same number of
valence electrons.
Example: Elements in Group 2A (2) have two (2)
valence electrons.
Be 1s22s2
Mg 1s22s22p63s2
Ca [Ar]4s2
Sr [Kr]5s2
Electron-Dot Symbols
An electron-dot symbol
§  indicates valence electrons
as dots around the symbol of
the element.
§  of Mg shows two valence
electrons as single dots on the
sides of the symbol Mg.
Mg
Mg
Mg
Mg
Mg
Periodic Table and
Valence Electrons
Writing Electron-Dot Symbols
The electron-dot symbols for
§  Groups 1A (1) to 4A (14) use single dots:
Na
Mg
Al
C
§  Groups 5A (15) to 7A (17) use pairs and single dots:
P
O
Cl
Groups and Electron-Dot Symbols
In a group, all the electron-dot symbols have the same
number of valence electrons (dots).
Example: Atoms of elements in Group 2A (2) each have
2 valence electrons.
Group 2A (2)
· Be ·
· Mg ·
· Ca ·
· Sr ·
· Ba ·
Atomic Size
Atomic size §  is described using the atomic radius. §  is the distance from the nucleus to the valence electrons. §  increases going down a group. §  decreases going across a period from lej to right. Atomic Radius
Ionization Energy
Ionization energy
§  is the energy it takes to remove a valence electron from
an atom in the gaseous state.
Na(g) + Energy (ionization)
Na+(g) + e–
§  decreases down a group, increasing across the periodic
table from left to right.
Ionization Energy and Valence Electrons
Ionization Energy
The ionization
energies of
§  metals are low.
§  nonmetals are high.
Learning Check
Identify the number of valence electrons for each of the
following:
1. 1s22s22p63s23p1
2. 1s22s22p63s2
3. 1s22s22p5