Download Group I Elements

AP Biology
Chapter 2
The chemical context of Life
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Chemical elements and Compounds
The atomic structure determines the behavior of elements.
Atoms combine chemically to form compounds and molecules.
A molecules biological function is related to its shape.
Chemical reactions break chemical bonds.
Organisms are composed of matter.
Life requires about 25 chemical elements
96% Living Matter
Carbon
Hydrogen
Oxygen
Nitrogen
4 % Remaining
Phosphorus
Sulfur
Calcium
Potassium
Trace elements are in organisms in minute quantities. Iron and iodine are
examples of these.
Atoms and Molecules
The nucleus of the atoms contains protons and neutrons.
ELECTRON
Mass = 9.11 x1031 kgms
Charge = -1
PROTON
Mass = 1.67 x1027 kgms
Charge = +1
NEUTRON
Mass = 1.67 x1027 kgms
Charge = 0
The nucleus of an atom has a high density compared to the overall atom.
The nucleus is very small compared to the overall atom.
The arrangement of the electrons determines the atoms chemical properties.
Atoms can have isotopes. Isotopes have the same number of protons and
electrons but a different number of neutrons.
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Atomic Number is the number of electrons and protons.
The total mass of the atom is Atomic Weight.
23
11
The element name is sodium
Na
23 is the mass number
11 is the atomic number
Molecules and Bonding
Forces that hold atoms together are chemical bonds.
Covalent bonds share electrons ---These include compounds such as water, alcohol, and most
organics.
Most compounds that are covalently bonded form within the “p”
block elements. Organic compounds are covalent in nature and
usually consist of carbon, hydrogen, oxygen, and nitrogen.
Ionic Bonds exchange electrons ----These include inorganic salts like NaCl or NaOH
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Binary compounds are formed by the action of metals reacting with
non-metals and mutually transferring electrons. This results in the
formation of cations and anions that mutually attract each other.
This attraction causes a change in energy resulting in the formation
of a binary or ionic compound. Energy is also a by-product.
Even though a solid compound is formed, for the electrons to
transfer a phase change takes place resulting in an energy change.
This energy is Latice Energy or the change in energy that takes
place when separated gaseous ions form together to make an ionic
solid.
The following reaction can happen:
M+(g) + X-(g)  MX(s) + ENERGY
Reactions can happen in other compounds.
Chemical Formula H 2O
Str
uct
ura
l
For
mu
la
Consider the ionic reaction: Na  Na   e  Sodium looses an electron
(OXIDATION)
Element
Consider the ionic reaction: Cl  e  Cl 
(REDUCTION)
Element
Cation or charged particle
Chlorine gains an electron
Anion or charged particle
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Consider Avogadro’s Hypothesis for chemical reactions
Polyatomic ions consists of several elements or atoms:
NH 4Cl
Ca3 PO4
AmmoniumChloride
Calcium Phosphate
morethantwo elements
Group I Elements
1) A +1 charge
2) Also calles the Alkali Metals.
3) React with water to form an hydroxide.
4) These elements do not occur in nature but in the form of
ores.
5) Ionize very easily.
Group II Elements
1) +2 Charge
2) Also called the Alkaline Earth Metals
3) React slowly with water to produce a metal oxide.
4) These elements also do not occur in nature but rather in
ores.
Transition Elements
1) Mostly alloy with other elements especially with steel to
make it tougher, ware resistant, and more durable.
2) Several of these elements have multi oxidation states.
3) The salts of these elements are very colorful.
4) Most of these elements conduct electricity.
Halogen Gases
1) These elements are all diatomic.
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2) These elements ionize very well especially with Group I
elements and somewhat with Group II elements.
3) All have a –1 charge.
Noble Gases
1) All are non-reactive.
2) All have a “0” charge.
3) Occur in nature as a mixture with other gases.
Lanthanides and Actinides
These elements are for the most part radioactive and used
for military purposes or medical applications.
Atomic Theory
Atomic orbitals:
Electrons can only occupy so-called atomic orbitals with well defined
energy levels corresponding to the principal quantum number, n. The
lowest level will have n = 1, the next n = 2, and so on.
The maximum number of electrons which can "fit" into these energy levels
is given by the formula 2n2, where n is the principal quantum number.
So, the first energy level will accommodate a maximum of 2 electrons (2 x
1 x 1), the second level 8 (2 x 2 x 2), the third 18 (2 x 3 x 3) and so on.
Within a main energy level with a specific value of n, there may be
subsidiary levels which are designated s, p, d and f in increasing order of
their energy values.
The diagram on the left shows the relative
energies of the atomic orbitals for principal
quantum numbers n = 1 to n = 4.
Note that the level with n = 1 is limited to a
single s orbital.
The level n = 2 has both s and p orbitals, while
at the level n = 3, we have d orbitals.
At the level n = 4, we have, in addition to s and
p orbitals, d and f orbitals, which are not
shown in the diagram.
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The Aufbau Process and Energy Levels
It is quite easy to establish this distribution for a given atom, by making
use of simple rules.
The best way to learn these rules is to see how they are applied in
practice, in the so-called AUFBAU (building-up) process, for atoms in the
ground state (that is, with all of its electrons in the lowest possible energy
levels).
Electrons must always enter the first
available orbital of lowest energy.
The first element, hydrogen, only has one
electron, and so this electron must enter the 1s
orbital.
The electronic configuration of hydrogen in the
ground state must therefore be: H 1s1
Pauli's exclusion principle must now be
applied - the next electron to enter the 1s
orbital must have a spin opposite to the
spin of the electron which is already there.
This completes the occupancy of the 1s orbital.
The electronic structure of helium is: He 1s2.
Consider the example
The next two electrons enter the 2s orbitals in the same way. This leads
us first to the element lithium, with 3 electrons (1s2 2s1) and then,
beryllium, with 4 electrons, and an electronic structure 1s2 2s2.
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The next element, boron, has 5 electrons.
The first four of these fill the 1s and 2s
levels, while the fifth enters the next level
of lowest energy, which is a 2p orbital. The
electronic structure of boron
is: 1s22s2 2p1.
We now apply Hund's rule: fill a set of
orbitals of equivalent energy (the 2p orbitals in
this case) in such a way that as many electrons
as possible remain unpaired.
The 6th electron enters a vacant 2p orbital rater
than pairing with an unpaired electron. This
element is carbon: 1s22s2 2p2.
In this way, the next electron enters the vacant 2p orbital, giving nitrogen
(1s22s2 2p3), oxygen (1s22s2 2p4), fluorine (1s22s2 2p5) and neon (1s22s2
2p6).
s-Orbitals
An orbital may be considered as a region in space to which electrons have
access. Heisenberg's uncertainty principle states that one cannot
simultaneously determine with certainty both the position and the
momentum of an electron. So, an orbital may be represented as a region
where there is a given probability of finding an electron at any given time.
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The diagram on the right shows how the
probability density, (P), of a 1s electron varies
with its distance (r) from the nucleus.
This so-called electron density of an electron in a
1s state is the same in all directions.
A 1s orbital may be considered as a spherical
distribution of negative charge, which becomes
more diffuse as the distance from the nucleus
increases.
One tries to depict this graphically, as shown
on the right. This is often called an electron
cloud.
2s orbitals are similar in shape, but the radius
of the electron cloud is larger.
p-Orbitals
p-Orbitals consist of electron clouds which look like
the diagram on the right, with a node at the nucleus of
the atom. There are three such orbitals, oriented at right
angles to one another, as shown in the diagram below:
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Combining all three p orbitals on the same set of
axes, as shown on the left, gives an idea (but
only a rough idea!) of the space available to an
electron in the p-state.
Remember that each orbital can only hold a maximum of 2 electrons
(which must have opposite spins). The complete set of the three orbitals
px, py and pz can therefore altogether accommodate a maximum of 6
electrons.
Ions:
Atoms are electrically neutral, but they may either lose or gain electrons,
becoming charged ions in the process. If an atom loses one or more
electrons, it will become positively charged and the resulting ion is called
a cation (see a table of common cations):
On the other hand, by gaining one or more electrons, an atom becomes
negatively charged, and is then called an anion (see a table of common
anions):
Let's see what happens to the electron configuration of a sodium atom
when it loses an electron to become the cation
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The change in electron configuration is described by
But 1s2 2s2 2p6 is the electron configuration of the inert gas neon (Ne),
which, like others in that group, shows remarkable stability.
What happens with chlorine? It gains an electron to form the chloride
anion
The change in electron configuration is described by
But 1s2 2s2 2p6 3s2 3p6 is the electron configuration of the element 18,
which is the inert gas argon (Ar). In general, metallic elements tend to lose
electrons to form a stable octet of electrons. The more reactive non-metals
such as oxygen and chlorine achieve this octet by gaining electrons.
Light energy as from the sun can excite electrons to form new
compounds. This concept is demonstrated in photosynthesis.
Some common Cations
Charge = +1
Charge = +2
Charge = +3
Hydrogen H+
Magnesium Mg2+
Aluminium Al3+
Lithium Li+
Calcium Ca2+
Iron(III) Fe3+
Sodium Na+
Barium Ba2+
Potassium K+
Zinc Zn2+
Silver Ag+
Iron(II) Fe2+
Copper(I) Cu+
Copper(II) Cu2+
Ammonium NH4+
Lead(II) Pb2+
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Some common anions:
Charge = -1
Charge = -2
Charge = -3
Fluoride FChloride ClBromide BrIodide IHydroxide OH-
Oxide O2-
Hydrogen sulphide(1) HS-
Sulphide S2-
Nitrite NO2Nitrate NO3Hydrogen carbonate(2)
HCO32-
Carbonate
CO22-
Hydrogen sulphite(3) HSO3-
Sulphite SO32-
Hydrogen sulphate(4) HSO4-
Sulphate SO42Phosphate
PO43-
(1)
Also known as "bisulphide" or "disulphide".
Also known as "bicarbonate".
(3)
Also known as "bisulphite".
(4)
Also known as "bisulphate".
(2)
Weak Chemical Bonds Play an Important Role in the Chemistry of
Life
Most chemical bonds in living organisms are covalent.
Some chemicals in a cell contact forming a weak chemical bond.
These weak bonds will respond to chemicals in another cell thus
forming a response. This response is similar to brain cells
responding to some stimulus.
Hydrogen bonding occurs when a hydrogen atom bonds covalently
to an electronegative atom that is also bonded to another
electronegative atom.
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The molecules which have this extra bonding are:
Note: The solid line represents a bond in the plane of the screen or
paper. Dotted bonds are going back into the screen or paper away from
you, and wedge-shaped ones are coming out towards you.
Notice that in each of these molecules:


The hydrogen is attached directly to one of the most
electronegative elements, causing the hydrogen to acquire a
significant amount of positive charge.
Each of the elements to which the hydrogen is attached is
not only significantly negative, but also has at least one
"active" lone pair.
Lone pairs at the 2-level have the electrons contained in a relatively
small volume of space which therefore has a high density of
negative charge. Lone pairs at higher levels are more diffuse and
not so attractive to positive things.
Another example of hydrogen bonding is water with ammonia. The
electronegative atoms are nitrogen and hydrogen. The water
orientated itself with the negative nitrogen atom forming a weak
covalent bond.
A non-polar molecule can have regions of “+” and “-“ causing weak
bonds to be temporally formed. These weak attractions are Vander
Waal forces.
Biological Function Related to Shape
The molecular shape of molecules (hybridization) determines they
respond and recognize another molecule and react with it.
Dalton’s Atomic Theory revolutionized chemistry by explaining
chemical properties in terms of small, indivisible pieces of matter
called atoms that are linked together to form polyatomic species
(both ions and molecules). As chemists explored the properties of
the polyatomic species, it became clear that they have size and
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shape and that shape is particularly important in explaining their
physical properties and why and how chemical reactions occur. A
summary of some of the commonly observed polyatomic shapes
(also known as structures or geometries) are shown in Table 1
arranged by the number of atoms around the "central atom." These
are idealized structures; real molecules seldom exhibit these
idealized shapes. However, the ideal shapes are good starting
points toward understanding how the spatial arrangement of atoms
in polyatomic species affect their properties and chemistry.
Number of Atoms
Around “Central
Atom”
Shape
1
Linear
2
Linear
Bent
Trigonal planar
3
Trigonal pyramidal
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T-shape
Tetrahedral
4
Square planar
Bisphenoid (see-saw)
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5
Trigonal bipyramidal
Square pyramidal
6
Octahedral
Chemical Reactions Make or Break Chemical Bonds
Chemical reactions demonstrate the conservation of mass and
energy.
In any chemical reaction chemical bonds are broken, new
substances are formed.
When the products formed equal the breakdown of the reactants
then chemical equilibrium happens.
Now that we know the how and why of chemical bonding, we can
look at some chemical reactions. Chemical reactions happen all
around us: when we light a match, start a car, eat dinner or walk
the dog. A chemical reaction is the pathway by which two
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substances bond together. In fact we have already discussed
several chemical reactions. One we have mentioned is the reaction
of hydrogen with oxygen to form water. To write the chemical
reaction you would place the reactants (the substances reacting) on
the left with an arrow pointing to the the products (the substances
being formed). Given this information, one might guess that the
reaction to form water is written:
H+O
H2O
However there are 2 problems with this chemical reaction. First,
because atoms like to have full valence shells, single H or O atoms
are rare (and unhappy) creatures. As we saw in the previous
lesson, both hydrogen and oxygen react with themselves to form
the molecules H2 and O2, respectively. These hydrogen and
oxygen molecules are much more common. Given this correction,
one might guess that the reaction looks like this:
H2 + O2
H2O
But we still have one problem. As written, this equation tells us that
1 hydrogen molecule (with 2 H atoms) reacts with 1 oxygen
molecule (with 2 O atoms) to form 1 water molecule (with 2 H
atoms and 1 O atom). In other words, we seem to have lost 1 O
atom along the way! To write a chemical reaction correctly, the
number of atoms on the left side of a chemical equation has to be
precisely balanced with the atoms on the right side of the
equation. How does this happen in our example? In actuality, the
O atom that we 'lost' reacts with a 2nd molecule of hydrogen to
form a second molecule of water. The reaction is therefore written:
2H2 + O2
2H2O
In the chemical reaction above, the number in front of the molecule
(called a coefficient) indicates how many molecules participate in
the reaction. A simulation of the reaction can be viewed by clicking
below (the atoms are represented as spheres in the animation: red
= hydrogen, blue = oxygen):
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