Download Chapter 2 – Atoms and Elements

Chapter 2 – Atoms and Elements
What is chemistry? It is often defined as “the study of matter”.
It answers the questions:
• “What is a substance made of?”
• “How was it made?”
• “How will it interact with other substances?”
e.g. The chemistry of beer
Beer is a homogeneous mixture consisting of water (_____),
ethanol (___________), carbon dioxide (_____) and a variety of
other substances responsible for its flavour.
Beer is made in a multi-step process:1
1. Barley mash is heated with water. This activates enzymes
in the barley that break the starch down to glucose.
2. The barley husks are filtered out of the resulting sugary
water (the “wort”) which is then boiled with hops to
impart flavour (by dissolving some of the more flavourful
molecules from the hops).
3. The hops are filtered out, and yeast is added for the
fermentation step in which they convert glucose (C6H12O6)
into carbon dioxide (CO2) and ethanol (CH3CH2OH):
C6H12O6(aq)
2 CO2(g) + 2 CH3CH2OH(aq)
4. After fermentation is complete, the yeast is filtered out.
The beer is then aged in tanks and filtered again before
packaging.
1
www.sleeman.ca, visited June 6, 2005
How does beer interact with other substances?
If certain bacteria get into the beer, their enzymes will
convert the ethanol in beer to acetic acid (CH3CO2H):
CH3CH2OH(aq) + O2(g)
CH3CO2H(aq) + H2O(l)
The interactions between beer and the human body are well
known (taste, inebriation, etc.) The taste is due to the
shapes of the flavour molecules and how they fit into
receptor molecules in our taste buds. Due to their structure,
the ethanol molecules travel easily through the human body
(they are soluble in both water and fat) until they reach
their target, the brain.
Chemistry is often termed “the central science”. It lies between
biology and physics, and chemical knowledge is also necessary
to truly understand many concepts in physics, biology,
medicine, geology, environmental science, etc.
In many ways, learning chemistry is like learning a language.
The chemist’s alphabet is the collection of elements – which are
combined to make compounds (“words”) – which are, in turn,
combined in reaction equations (“sentences”). Just as there are
rules for writing sentences that make sense, we will see that
there are rules for writing reaction equations that make sense.
Atomic Theory
Ancient Greek philosophers proposed that all matter consisted of
some combination of four elements: air, earth, fire, water.
Democritus (~460-370 B.C.) disagreed, proposing that all matter
could be repeatedly subdivided until an indivisible particle was
reached. He called this the atom (Greek: a = not; tomos = cut).
In 1785, Antoine Lavoisier (1743-1794) is credited with
discovery of the law of conservation of mass: Mass is neither
created nor destroyed in a chemical change.
In 1794, Joseph Proust (1754-1826) demonstrated the law of
definite proportions (aka law of constant composition): In a
given chemical compound, the proportions by mass of the
elements that compose it are fixed, independent of the origin of
the compound or its mode of preparation.
In 1808, John Dalton (1766-1844) based his atomic theory of
matter on the ideas of Democritus, Lavoisier and Proust.
1. All matter consists of solid and indivisible atoms.
2. All of the atoms of a given chemical element are identical
in mass and in all other properties.
3. Different elements have different kinds of atoms; these
atoms differ in mass from element to element.
4. Atoms are indestructible and retain their identity in all
chemical reactions.
5. The formation of a compound from its elements occurs
through the combination of atoms of unlike elements in
small whole-number ratios.
While the essence of this theory has withstood the test of time,
most of the postulates have since been modified:
1. Atoms can be further divided into subatomic particles.
2. Different isotopes of an element have different masses.
(e.g. Carbon-12, carbon-13 and carbon-14 have masses of
12, 13.003 and 14.003 u respectively.)
3. Still true, but some have very similar masses. (e.g.
Nitrogen-14 and carbon-14 both have masses of 14.003 u.)
4. In nuclear reactions, atoms do not retain their identity.
5. True; however, Dalton was unaware that not all elements
are made up of single atoms. e.g. Nitrogen does not exist
as free atoms; rather, two nitrogen atoms (N) bond to make
one nitrogen molecule (N2).
A molecule is a grouping of two or more atoms bonded together
by strong attractive forces. A molecule is a discrete entity.
Cl
Cl
Na
Na
O
H
H
H
Na
H
Cl
Cl
Na
Cl
Na
Na
Cl
molecules
Cl
Na
Na
Cl
Cl
Na
NaCl does not exist as discrete molecules
Elements can exist as:
• atoms (e.g. He, Ne)
• molecules (e.g. H2, N2, O2, F2, Cl2, Br2, I2, P4, S8)
• infinite materials (e.g. most metals and metalloids,
carbon as diamond or graphite).
Compounds can exist as:
• molecules (e.g. H2O, CO, CO2)
• infinite materials (e.g. SiO2 can be either quartz or sand)
• simple ions (e.g. NaCl consists of Na+ and Cl-)
• complex ions (e.g. CuSO4 consists of Cu2+ and SO42-).
Compounds and elements can be described by chemical and
physical properties:
A _____________________ is one that can be observed without
changing a compound/element into another compound/element.
A _____________________ is one that can only be observed by
changing a compound/element into another compound/element.
Define the following properties as chemical or physical:
1. Water has a density of 1.000 g/mL at 4˚C.
2. Sodium reacts with water to produce sodium hydroxide and
hydrogen gas.
3. Gasoline burns in the presence of oxygen to give carbon
dioxide and water vapour.
4. Solid iodine sublimes to give dark purple iodine vapour.
…back to Atomic Theory
In 1897, Sir John Joseph Thomson (1856-1940)
discovered the electron.2 He also proposed that
atoms contained electrons. He reasoned that, since
electrons have negative charge and atoms are
neutral, atoms must also contain some constituent
with positive charge. His proposal was the ‘plum
pudding’ model of the atom in which electrons
were embedded in a positively charged sphere.
2
previously proposed and named by physicist G. B. Stoney in 1874
+
+
-
-
+
-
+
-
+
+
+
+
-
'plum pudding' model
of the atom
While Thomson was able to measure the charge-to-mass ratio
for the electron, Robert Andrews Millikan (1868-1953) was the
first to measure the charge independently (-1.602176 × 10-19 C).
This also allowed calculation of the mass (9.109383 × 10-28 g).
As a matter of convenience, chemists refer to charge relative to
the charge of an electron rather than in Coulombs (C). By this
method, the electron has a charge of -1.
In 1911, Ernest Rutherford (1871-1937) proposed a new model
of the atom. He found that most alpha
nucleus (charge of +78)
particles fired at a thin gold foil passed
78 electrons
through the foil as though traveling through
empty space; however, a few were deflected
by large angles. This was inconsistent with
Thomson’s model. Rutherford’s ‘nuclear’
model consisted of a tiny positively charged
nucleus containing most of the atom’s mass. nuclear model of a gold atom
In 1919, Rutherford reported discovery of the proton,
demonstrating that the nucleus of a hydrogen atom was the
fundamental unit of positive charge and proposing that the
nuclei of heavier atoms consisted of protons and neutral
particles of similar mass. These neutral particles, ___________,
were discovered by James Chadwick (1891-1974) in 1932.
Thus, atoms consist of three types of subatomic particles:
mass
charge location
outer
electron 9.109383 × 10-28 g 0.0005485799 u
-1
region
-24
proton 1.672622 × 10 g
1.007276 u
+1
nucleus
-24
neutron 1.674927 × 10 g
1.008665 u
0
nucleus
Defining an Element
Given the extremely small masses of atoms and subatomic
particles, chemists invented a new unit of mass. The atomic
mass unit (u) is defined as one twelfth of the mass of a carbon
atom containing six protons, six neutrons and six electrons:
1 u = 1.661 × 10-24 g
As such, the mass of an atom in u will be approximately equal to
the combined number of protons and neutrons it contains.
Every atom has an atomic number and a mass number:
mass number
symbol
atomic number
12
6
C
Atomic number (Z) = # protons
Mass number (A) = # protons + # neutrons
The charge of an atom can be found by comparing the number
of protons and electrons:
charge = # protons - # electrons
A neutral atom has an equal number of protons and electrons.
The number of protons defines what element an atom belongs
to. If the number of protons changes, it is not the same element.
As such, writing atomic number is optional because the element
symbol tells us what it must be. Mass numbers are not optional.
How many protons, neutrons and electrons are in the neutral
element with Z = 26 and A = 56? What element is this?
Isotopes
Different atoms of the same element can have different numbers
of neutrons. They will have the same ____________ number
but different _________ numbers.
e.g.
1
2
1
1
12
13
H
6
C
6
H
3
C
14
1
6
H
C
These are called isotopes of the element.
Only a few elements have just one naturally occurring isotope
(e.g. 19F, 31P). Most elements occur as mixtures of several
isotopes. Chemists normally treat these elements as consisting
of “averaged” atoms with a “averaged” masses.
Atomic mass (as shown on the periodic table) is the weighted
average of all naturally occurring isotopes of an element. It
factors in the mass and percent abundance of each isotope where
% abundance =
# atoms of isotope
× 100%
total # atoms of element
atomic mass =
%abund. iso.#1
%abund. iso.#2
(mass iso.#2)
(mass iso.#1) +
100%
100%
e.g. Chlorine has two naturally occurring isotopes:
35
37
17
17
Cl
Cl
75.8%
24.2%
34.97 u
36.97 u
Calculate the average atomic mass of chlorine.
e.g. Gallium has two naturally occurring isotopes and an average
atomic mass of 69.723 u:
69
71
31
31
Ga
Ga
68.926 u
70.925 u
Calculate the percent abundance of each isotope of gallium.
Thus, we can calculate the mass of an atom. By adding atomic
masses, we can also calculate the mass of a molecule (or
formula unit of an infinite material).
e.g. Sulphur has an atomic mass of 32.066 u.
Oxygen has an atomic mass of 15.999 u.
Calculate the molecular mass of SO3.
In the lab, however, chemists don’t often work with single atoms
or molecules. It is far more common to work with quantities in
the 1 mg to 1 kg range. Recall that 1 u = 1.661 × 10-24 g. Thus,
1 g = 6.022 × 1023 u. This means that 1 gram contains a *lot* of
atoms or molecules. Since the human brain has trouble
comprehending numbers this large, another unit was created to
make such quantities easier to discuss:
1 mole = 6.022 × 1023
Thus, 1 g = 1 mole µ
Or, 1 µ = 1 g/mole
This ‘magic’ number (6.022 × 1023) is called _______________
_______________ in honour of Amedeo Avogadro (1776-1856)
who was the first to propose that such a number could exist.
For convenience, the word ‘mole’ is often abbreviated to ‘mol’
when used as a unit. e.g. The mass of 12C is 12 g/mol.
The key to working with moles is to remember that 1 mole is
always equal to 6.022 × 1023 of whatever type of particle you are
discussing.
e.g. There are __________ molecules of CO2 in 1 mole of CO2.
There are __________ atoms of C in 1 mole of CO2.
There are __________ atoms of O in 1 mole of CO2.
How many moles of C2H6O contain 5.0 x 1024 atoms of H?
How many molecules are there in 392.3 g sulfuric acid (H2SO4)?
What is the mass of a sample of hydrochloric acid (HCl)
containing 2.01 × 1024 atoms of hydrogen?
What is the mass of a sample of nitric acid (HNO3) containing
3.011 × 1022 atoms of oxygen?
The Elements
Currently, ____ elements are known. Of these, ____ occur
naturally while the other ____ have been made in laboratories.
The Gases
____ elements are gases at room temperature
The “noble gases” or “inert gases” are:
Helium, Neon, Argon, Krypton, Xenon, Radon
Air is mostly Nitrogen (___) and Oxygen (___)
(~80% N2 + ~20% O2)
The other gases are:
Fluorine (symbol ___, formula ____, highly reactive)
Chlorine (symbol ___, formula ____, reactive)
Hydrogen (symbol __, formula ___, flammable)
The Liquids
____ elements are liquids at room temperature:
Bromine (symbol ____, formula ____, corrosive)
Mercury (symbol ____, formula ____, poisonous)
The Solids
The remaining elements are solids at room temperature.
Chemists have a special tool for recognizing which elements are
similar and which are different. This tool is known as the
________________________.
The Periodic Table
In 1869, Dmitri Mendeleev (1834-1907) noticed that certain
elements exhibited similar behaviour – most notably, the ratios
with which they formed molecules with hydrogen and with
oxygen. By arranging the elements in order of increasing mass
and such that similar elements formed columns, he developed
the first periodic table:
This periodic table was incomplete – all of the noble gases are
missing, but it was remarkably accurate in other respects. If
there appeared to be a ‘missing’ element, Mendeleev left a blank
space, assuming that it would be discovered at a later date. He
was proven correct with the discoveries of _________ (69.7 u)
in 1875 and _______________ (72.6 u) in 1886.
In 1913, H.G.J. Moseley (1887-1915) noted that the periodic
table would be more descriptive if the elements were listed in
order of increasing _____________ rather than increasing mass.
This led to the modern periodic table and law of periodicity:
“The properties of the elements are periodic functions of atomic
number.” We will see why this is after looking at some of the
properties of the elements.
In the modern periodic table, a column is known as a
________________ and a row is known as a _______________.
Elements in the same period have similar ___________ but very
different ________________. Elements in the same group have
similar _________________ but very different ____________.
Despite having similar chemical properties, elements in the
same group sometimes have quite different physical properties.
Why might this be?
main
groups
The periodic table can also be divided
to give different classes of elements.
The upper right portion of the periodic
table consists of nonmetals while the
lower left portion consists of metals
and a few elements at the border are
metalloids.
main
groups
transition
groups
A Quick Overview of the Elements by Group
Hydrogen (H)
• shares similarities with elements in groups 1 and 17 but
doesn’t fully belong to either group
• diatomic gas (H2) that is fairly unreactive
• reacts explosively with oxygen to make water (H2O) if
enough energy is supplied
• three isotopes: 1H is the most common (and the only
element with no neutrons), 2H is in heavy water used to
cool nuclear reactors; 3H is radioactive
Group 1: Alkali Metals (Li, Na, K, Rb, Cs, Fr)
• aka Group 1A
• soft metals that react strongly with water to give XOH
and with oxygen to give X2O (reactivity increases with
atomic mass)
• do not exist in pure form in nature due to high reactivity
• readily lose 1 electron to make cations with +1 charge
Group 2: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)
• aka Group 2A
• most are metals that react with water to give X(OH)2
and with oxygen to give XO (reactivity increases with
atomic mass; beryllium does not react with water)
• do not exist in pure form in nature due to high reactivity
(though react less strongly with water than the Group 1
element in the same period)
• beryllium is highly toxic
• readily lose 2 electrons to make cations with +2 charge
Groups 3-12: Transition Metals
• aka Groups 1B-8B (but not in order!)
• Group 3 metals readily lose 3 electrons to make cations
with +3 charge; otherwise, they act like Group 2 metals
• Groups 4-11 are the ‘true’ transition metals in that they
lose electrons to form coloured compounds in which the
metal atom has a positive charge
• Groups 11 and 12 mimic the behaviour of Groups 1 and
2 respectively but are less reactive
• the metals in the middle of the transition groups are
hardest (especially tungsten (W) in Group 6)
• the metals at the edges of the transition groups (i.e.
Groups 3 and 12) are the softest
• mercury is the only liquid metal; the rest are solids
• copper and gold are the only coloured metals
• silver is the best conductor of electricity
• gold is the most malleable metal
• readily lose electrons to make cations
Group 13 (B, Al, Ga, In, Tl)
• aka Group 3A
• most are metals; boron is a metalloid
• all are solids; however, gallium has a very low melting
point (29.76 ˚C)
• aluminum is an industrially important lightweight metal
and the third most abundant element in the earth’s crust
• compounds containing indium and thallium are toxic
• form compounds in a 1:3 ratio with halogens (e.g. BCl3)
• lose 3 electrons to make cations with +3 charge
Group 14 (C, Si, Ge, Sn, Pb)
• aka Group 4A
• carbon is a nonmetal; silicon and germanium are
metalloids; tin and lead are metals
• carbon exists in several different allotropes (graphite,
diamonds, fullerenes), and compounds containing
carbon form the basis for life
• silicon is the second most abundant element in the
earth’s crust; it does not naturally occur in pure form but
as silicates (compounds made of silicon and oxygen)
which form rocks, sand, glass, etc.
• form compounds in a 1:4 ratio with halogens (e.g. CCl4)
Group 15: Pnictogens (N, P, As, Sb, Bi)
• aka Group 5A
• nitrogen and phosphorus are nonmetals; arsenic and
antimony are metalloids; bismuth is somewhat metallic
• nitrogen is a highly stable diatomic gas (N2) and the
most abundant element in the atmosphere
• phosphorus exists in three allotropes (white phosphorus
is P4; red and black phosphorus are polymer used in
match heads)
• form compounds in a 1:3 ratio with hydrogen (e.g. NH3)
Group 16: Chalcogens (O, S, Se, Te, Po)
• aka Group 6A
• oxygen, sulphur and selenium are nonmetals; tellurium
is a metalloid; polonium is a metal
• oxygen is the most abundant element in the earth’s crust
& the second most abundant element in the atmosphere
• oxygen exists in two allotropes, O2 (oxygen) and O3
(ozone); both are reactive gases
• sulphur exists in many allotropes (S2, S6, S8, etc.)
• except for oxygen, chalcogens tend to make compounds
with unpleasant odours (which get worse as atomic mass
increases)
• form compounds in a 1:2 ratio with hydrogen (e.g. H2O)
• gain 2 electrons to make anions with -2 charge
Group 17: Halogens (F, Cl, Br, I, At)
• aka Group 7A
• nonmetals that exist as diatomic molecules (except for
astatine which is too unstable to study)
• fluorine and chlorine are gases; bromine is a liquid;
iodine is a solid
• colourful (F2 is pale yellow; Cl2 is yellow-green; Br2 is
red-brown; I2 is dark purple)
• form compounds in a 1:1 ratio with hydrogen (e.g. HF)
• gain 1 electron to make anions with -1 charge
Group 18: Noble Gases (He, Ne, Ar, Kr, Xe, Rn)
• aka Group 8A, inert gases or rare gases
• unreactive gaseous nonmetals (a few compounds have
been made containing Xe and Kr but they are rare)
• helium is used to fill blimps, etc. due to low density
• 1% of the atmosphere is argon, and argon is often used
as an inert atmosphere in labs because it is denser than
air and unreactive
Important Concepts from Chapter 2
• chemical vs. physical properties
• law of conservation of mass
• Dalton’s atomic theory of matter
• Thomson and Rutherford models of the atom
• subatomic particles (protons, neutrons, electrons)
• atoms, molecules, infinite materials, simple and complex ions
• periodic table (groups and periods)
• elements (names and symbols)
• atomic number and mass number
• isotopes
• calculating average atomic mass and percent abundance
• Avogadro’s number and the mole