Download INTRO TO CHEMISTRY WEEK: SEPTEMBER 14 – 18, 2015

LS50 2015
INTRO TO CHEMISTRY WEEK: SEPTEMBER 14 – 18, 2015
Learning goals
By the end of this week, you should understand the following:
• The key findings and interpretations of some critical experiments that led to our current
understanding of atomic structure and electron configurations (Lecture 08)
• How electron configurations are related to chemical bonds and interactions (Lecture 09)
• How chemical bonds are related to assembly and structure of biological molecules (Lectures 0910)
• How molecular assembly and structure are related to molecular function (Lectures 11-12)
• How knowing molecular function facilitates prediction of the evolution of biological molecules
(Lecture 12)
Lecture 09 – Elements and periodic trends
Learning goals
By the end of this lecture, you should be able to:
• Know how to figure out the number of valence electrons for an atom
• Understand electronegativity; use this to predict formation of compounds in general terms
• Describe and explain some basic properties of an element based on periodic trends
• State the octet rule
o name and define the types of bonds that atoms can engage in to satisfy it
o recognize the three major types of violations of this rule
• Draw Lewis dot structures of atoms, ionic compounds, and covalent compounds
Lecture Outline
Periodic table
• note that atomic number is not the same as atomic weight:
o atomic number = number of protons (this is the order the elements are arranged in, and
is constant regardless of the number of neutrons)
o atomic weight = weighted average of the atomic mass numbers of all isotopes of an
element
o atomic mass number = (atomic number) + (number of neutrons)
= (number of protons) + (number of neutrons)
• atoms can have more than one
o ionic form (with positive (cation) or negative (anion) charge)
§ charge is determined by the relative number of protons and electrons:
§ charge = (number of protons) – (number of electrons)
o isotope, i.e. have a different number of neutrons
§ isotopes have the same atomic number but different atomic mass numbers and
different numbers of neutrons
• the elements break down into some broad categories (we will refine these later):
o metals
§ tend to lose electrons to become cations
§ tend to become isoelectric with a noble gas (this means that they achieve the
same number of electrons as a noble gas)
§ tend to be shiny (have luster), malleable, ductile, good conductors of heat and
electricity
§ six main types: alkali metals, alkaline earth metals, lanthanoids, actinoids,
transition metals and post-transition metals
o non-metals
§ tend to gain electrons to become anions
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o
§ tend to become isoelectric with a noble gas
§ tend to be dull, brittle, poor conductors of heat and electricity
§ three main types: noble gases, halogens, other non-metals
metalloids
§ can gain or lose electrons: no good way to predict what these will do, although Al
tends to be mostly metallic and B tends to be mostly nonmetallic
§ they border the amphoteric line on the periodic table
•
Valence electrons
o These are the electrons that can participate in forming chemical bonds!
o You can use the periodic table to help you figure out how many valence electrons an
atom has (see Table)
o Noble gases don’t have any valence electrons because their outermost shells are totally
full! That’s why they are inert
o Atoms can do different things with their valence electrons:
§ Contribute them to a shared partnership with another atom à covalent bond
§ Totally donate them or steal them from another atom à ionic bond
§ Nothing à no bond
o The valence electron number and position determine which one of these things will
happen = the chemical properties of the element
•
The octet rule
o (it is an empirical observation that) the atoms in the main group elements bond so that
each atom has 8 electrons (or 2 electrons if it is H or He) in its outermost shell
o can be satisfied by covalent bonds or ionic bonds
§ ionic compounds:
• form when the difference in electronegativity between two atoms is large
• an electron is totally transferred from one atom to another (the one with
higher electronegativity: further to the right of the periodic table), so that
they both satisfy the octet rule
§ covalent compounds:
• form when the difference in electronegativity between two atoms is
small(er)
• an electron is shared between two atoms, so that they both satisfy the
octet rule
•
Lewis dot structures of atoms
o A way to describe valence electrons in atoms, and bonding in polyatomic systems
o One electron pair = one single bond (between two atoms)
o Follow these steps to draw the Lewis dot structure for atoms:
§ Draw the atomic symbol
§ Represent each valence electron as a dot around it
o The location of the dots is not important; they will get moved around when assigning
electrons to bonds in compounds
o When drawing ions with partially filled d or f subshells (e.g. transition metals), typically we
don’t include those subshells in the drawing
o For cations or anions that might have up to 8 electrons in their outer shell, typically we
write the original valence shell configuration (the shell configuration pre-ionization)
•
Periodic trends
o The elements in a given group (column) have similar properties because their outer
electron configurations are similar to each other
o There are also trends across a period (row)
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o
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o
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As long as we arrange elements according to atomic number (assuming neutral atoms,
so that # protons = # electrons), the physical and chemical properties of the elements
vary periodically
Four major trends that help us understand how atoms behave in chemical bonds:
§ Atomic radius (and ionic radius)
• Decreases across a period from L to R
• Increases down a group
§ Ionization energy
• This is the minimum energy needed to remove the highest energy
electron from an atom (creating a cation)
• Increases across a period from L to R (gets harder to remove electrons L
to R across a period)
• Decreases down a group (gets easier to remove electrons going down a
group)
§ Electronegativity
• This is the net ability of an atom to take an electron from another atom
(i.e. be an oxidizing agent)
• Increases across a period from L to R (the elements get better at stealing
electrons)
• Decreases down a group (the elements get worse at stealing electrons)
§ Electron Affinity
• This is the energy change required to add an electron to a neutral atom
(creating an anion)
• Increases from bottom left to top right (gets harder to add electrons in
this direction)
A bit more detail about characteristics of some of the groups:
Alkali metals (Li, Na, K, Rb, Cs, Fr)
§ Group IA = Group 1 à one valence electron (in s subshell)
§ Have the largest atomic radius of their period
§ React violently with water à H2
§ Easily ionize to 1+ by losing an electron from s à empty s
§ Oxidize (lose electrons) in air
Alkali earth metals (Be, Mg, Ca, Sr, Ba, Ra)
§ Group IIA = Group 2 à two valence electrons (in s subshell)
§ Easily ionize to 2+ by losing 2 electrons from s à empty s
§ Close (filled) subshell
§ React with water à H2
Transition metals (Sc-Zn, Y-Cd, Hf-Hg, Rf-Cn)
§ Groups 3-12: multiple valence electrons in d subshells
§ May have more than one oxidation group
§ Reactive with acids
§ Make coloured compounds
§ Some groups have some some cool trivial names! See Table
Post-Transition metals (Al, Ga, In, Tl, Uut)
§ Group IIIA = Group 3A à Three valence electrons (in p subshell)
§ Can have several oxidation states, but commonly 3+
Group IVA = Group 4A (C, Si, Ge, Sn, Pb, Fl)
§ à 4 valence electrons (in p subshell)
§ Includes 1 nonmetal (C), 2 metalloids (Si, Ge) and 3 post-transition metals (Sn,
Pb, Fl)
§ Oxidation numbers +4 to -4
§ They form most covalent compounds
Group VA = Group 5A (N, P, As, Sb, Bi Uup)
§ à five valence electrons (in p subshell)
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Include 2 post-transition metals (Uup, Bi), 2 metalloids (As, Sb) and 2 nonmetals
(N, P)
§ Usually form anions -1, -2 or -3
Group VIA = Group 6A (O, S, Se, Te, Po, Lv)
§ à six valence electrons (in p subshell)
§ Include 3 nonmetals (O, S, Se), 2 metalloids (Te, Po) and 1 post-transition metal
(Lv)
Halogens (F, Cl, Br, I, At, Uus)
§ Group VIIA = Group 7A à seven valence electrons (full p subshell, one unpaired
in s subshell)
§ Form monoatomic gases
Noble gases (He, Ne, Ar, Kr, Xe, Rn, Uuo)
§ Group VIIIA – Group 8A à eight valence electrons (n level orbitals all filled)
§ Closed shell
§ Form monoatomic gases
Lanthanoids (La to Lu) and Actinoids (Ac to Lr)
§ Have a variable number of valence electrons in the f subshell
§
o
o
o
o
•
Lewis dot structures of ionic compounds
• Since electrons are not shared between atoms to achieve the octet rule in this type of bond,
just draw the structure appropriate for each ion following the electron transfer
• One of the atoms involved might end up with zero dots, and the other with eight dots
• If necessary indicate the charge in the top right hand corner of the atom
• The strength of the bond depends on the charge magnitude (greater charge à stronger
bond) and the ion size (smaller ion à stronger bond)
o Bond strength is called lattice energy, measured in kJ/mol
•
Lewis dot structures of covalent compounds
o Single bond = one electron pair shared = bond order 1
o Double bond – two electron pairs shared = bond order 2
o Triple bond = three electron pairs shared = bond order 3
o Quadruple bond = four electron pairs shared = bond order 4
o Fractional bond orders are also possible: see “Resonance” below)
o Sometimes the octet rule is violated! There are three main examples of this:
§ Sub-octet systems = octet-deficient molecules
• B and Be are prone to this: they can form stable molecules that don’t
obey the octet rule
• e.g. BF3, BeCl2, BCl2
§ Valence shell expansion
• This tends to happen for third period elements (Na, Mg, Al, Si, P, S. Cl)
• Their d orbitals are energetically close enough to be available to
participate in bonding
• >8 electrons can surround one of these elements in a molecule and it
can still form a stable bond
• e.g. ClF3, PCl5
§ Odd number of valence electrons = unpaired valence electrons
• e.g. ClO2
• these are free radicals: very reactive because of their unpaired electron
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Table. Characteristics of Groups (columns) of the Periodic Table
Example
Element
(main
group
only)
CAS
Group #
(older)
IUPAC
Group #
(modern)
# valence
electrons
Location
of
valence
electrons
Element
family
name
Trivial
name
H
Be
IA
IIA
IIIB
IVB
VB
VIB
VIIB
VIII
VIII
VIII
IB
1
2
3
4
5
6
7
8
9
10
11
1
2
Beryllium
Alkali
metals
Alkaline
earth
metals
C
N
O
F
He
IIB
IIIA
IVA
VA
VIA
VIIA
VIIIA
12
13
14
15
16
17
18
3
4
5
6
7
8 (0)
Transition metals: No good way to predict # valence electrons for these.
Outermost shells
Lithium
B
can be in their inner shells as well as outer
Scandium
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Titanium
Vanadium
Chromium
Manganese
Iron
Cobalt
Outermost shells
Nickel
Copper
Zinc
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Helium
Coinage
metals
Volatile
metals
Icosagens
Crystallogens
Pnictogens
Chalcogens
Halogens
Noble
gases