Download Lecture 1 – Matter, Atomic Structure

Chemistry
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Chemistry: A Science for the 21st Century
• Health and Medicine
• Sanitation systems
• Surgery with anesthesia
• Vaccines and antibiotics
• Gene therapy
•Energy and the Environment
• Fossil fuels
• Solar energy
• Nuclear energy
2
Chemistry: A Science for the 21st Century
• Materials and Technology
• Polymers, ceramics, liquid crystals
• Room-temperature superconductors?
• Molecular computing?
• Food and Agriculture
• Genetically modified crops
• “Natural” pesticides
• Specialized fertilizers
3
The Study of Chemistry
Macroscopic
Microscopic
4
The scientific method is a systematic
approach to research.
A hypothesis is a tentative explanation for a
set of observations.
tested
modified
5
A law is a concise statement of a relationship
between phenomena that is always the same
under the same conditions.
Force = mass x acceleration
A theory is a unifying principle that explains
a body of facts and/or those laws that are
based on them.
Atomic Theory
6
Chemistry is the study of matter and the
changes it undergoes.
Matter is anything that occupies space and
has mass.
A substance is a form of matter that has a
definite composition and distinct properties.
liquid nitrogen
gold ingots
silicon crystals
7
A mixture is a combination of two or more substances
in which the substances retain their distinct identities.
1. Homogenous mixture – composition of the
mixture is the same throughout
soft drink, milk, solder
2. Heterogeneous mixture – composition is not
uniform throughout
cement,
iron filings in sand
8
Physical means can be used to separate a mixture
into its pure components.
magnet
distillation
9
An element is a substance that cannot be
separated into simpler substances by chemical
means.
•114 elements have been identified
• 82 elements occur naturally on Earth
gold, aluminum, lead, oxygen, carbon, sulfur
• 32 elements have been created by scientists
technetium, americium, seaborgium
10
Replace with Table 1.1 from 7e page 6
11
A compound is a substance composed of atoms
of two or more elements chemically united in fixed
proportions.
Compounds can only be separated into their
pure components (elements) by chemical
means.
lithium fluoride
quartz
dry ice – carbon dioxide
12
Classification of Matter
Replace with Table 1.5 from 7e page 7
13
A Comparison: The Three States of Matter
14
The Three States of Matter: Effect of a Hot
Poker on a Block of Ice
15
Types of Changes
A physical change does not alter the composition
or identity of a substance.
sugar dissolving
ice melting
in water
A chemical change alters the composition or
identity of the substance(s) involved.
hydrogen burns in
air to form water
16
Extensive and Intensive Properties
An extensive property of a material depends upon
how much matter is being considered.
• mass
• length
• volume
An intensive property of a material does not
depend upon how much matter is being
considered.
• density
• temperature
• color
17
Matter - anything that occupies space and has mass
mass – measure of the quantity of matter
SI unit of mass is the kilogram (kg)
1 kg = 1000 g = 1 x 103 g
weight – force that gravity exerts on an object
weight = c x mass
A 1 kg bar will weigh
on earth, c = 1.0
1 kg on earth
on moon, c ~ 0.1
0.1 kg on moon
18
Atoms, Molecules, and Ions
Dalton’s Atomic Theory (1808)
1. Elements are composed of extremely small particles
called atoms.
2. All atoms of a given element are identical, having the
same size, mass and chemical properties. The atoms of
one element are different from the atoms of all other
elements.
3. Compounds are composed of atoms of more than one
element. In any compound, the ratio of the numbers of
atoms of any two of the elements present is either an
integer or a simple fraction.
4. A chemical reaction involves only the separation,
combination, or rearrangement of atoms; it does not
result in their creation or destruction.
20
Dalton’s Atomic Theory
Law of Multiple Proportions
21
16 X
+
8Y
8 X2Y
Law of Conservation of Mass
22
Cathode Ray Tube
J.J. Thomson, measured mass/charge of e(1906 Nobel Prize in Physics)
23
Cathode Ray Tube
24
Millikan’s Experiment
Measured mass of e(1923 Nobel Prize in Physics)
e- charge = -1.60 x 10-19 C
Thomson’s charge/mass of e- = -1.76 x 108 C/g
e- mass = 9.10 x 10-28 g
25
Types of Radioactivity
(uranium compound)
26
Thomson’s Model
27
Rutherford and the Nuclear Atom
• Ernest Rutherford directed Hans Geiger and
Ernst Marsden’s experiment in 1910.
– - particle scattering from thin Au foils
– Gave us the basic picture of the atom’s structure.
28
Rutherford’s Experiment
(1908 Nobel Prize in Chemistry)
 particle velocity ~ 1.4 x 107 m/s
(~5% speed of light)
1. atoms positive charge is concentrated in the nucleus
2. proton (p) has opposite (+) charge of electron (-)
3. mass of p is 1840 x mass of e- (1.67 x 10-24 g)
29
Rutherford and the Nuclear Atom
• In 1912 Rutherford decoded the -particle
scattering information.
– Explanation involved a nuclear atom with electrons
surrounding the nucleus .
30
Rutherford and the Nuclear Atom
Rutherford’s major conclusions from the -particle
scattering experiment
1. The atom is mostly empty space.
2. It contains a very small, dense center called the
nucleus.
3. Nearly all of the atom’s mass is in the nucleus.
4. The nuclear diameter is 1/10,000 to 1/100,000 times
less than atom’s radius.
31
Rutherford and the Nuclear Atom
• Because the atom’s mass is contained in such
a small volume:
– The nuclear density is ~1015g/mL.
– This is equivalent to ~3.72 x 109 tons/in3.
– Density inside the nucleus is almost the same as a
neutron star’s density.
32
Rutherford’s Model of
the Atom
atomic radius ~ 100 pm = 1 x 10-10 m
nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
“If the atom is the Houston
Astrodome, then the nucleus is a
marble on the 50-yard line.”
33
Chadwick’s Experiment (1932)
(1935 Noble Prize in Physics)
H atoms: 1 p; He atoms: 2 p
mass He/mass H should = 2
measured mass He/mass H = 4
 + 9Be
1n
+ 12C + energy
neutron (n) is neutral (charge = 0)
n mass ~ p mass = 1.67 x 10-24 g
34
mass p ≈ mass n ≈ 1840 x mass e35
Atomic Number, Mass Number, and Isotopes
Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different
numbers of neutrons in their nuclei
Mass Number
A
ZX
Atomic Number
1
1H
235
92
2
1H
U
Element Symbol
(D)
238
92
3
1H
U
(T)
36
The Isotopes of Hydrogen
37
A molecule is an aggregate of two or more atoms in a
definite arrangement held together by chemical forces.
H2
H2O
NH3
CH4
A diatomic molecule contains only two atoms:
H2, N2, O2, Br2, HCl, CO
diatomic elements
A polyatomic molecule contains more than two atoms:
O3, H2O, NH3, CH4
38
An ion is an atom, or group of atoms, that has a net
positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
Cl-
17 protons
18 electrons
39
A monatomic ion contains only one atom:
Na+, Cl-, Ca2+, O2-, Al3+, N3-
A polyatomic ion contains more than one atom:
OH-, CN-, NH4+, NO3-
40
2.1
Give the number of protons, neutrons, and electrons in each of
the following species:
(a)
(b)
(c)
(d) carbon-14
2.1
Strategy Recall that the superscript denotes the mass number (A)
and the subscript denotes the atomic number (Z).
Mass number is always greater than atomic number. (The only
exception is H, where the mass number is equal to the atomic
number.)
In a case where no subscript is shown, as in parts (c) and (d), the
atomic number can be deduced from the element symbol or name.
To determine the number of electrons, remember that because atoms
are electrically neutral, the number of electrons is equal to the
number of protons.
2.1
Solution
(a)
The atomic number is 11, so there are 11 protons.
The mass number is 20, so the number of neutrons is
20 − 11 = 9. The number of electrons is the same as the
number of protons; that is, 11.
(b)
The atomic number is the same as that in (a), or 11.
The mass number is 22, so the number of neutrons is
22 − 11 = 11. The number of electrons is 11. Note that the
species in (a) and (b) are chemically similar isotopes of
sodium.
2.1
(c)
The atomic number of O (oxygen) is 8, so there are
8 protons. The mass number is 17, so there are 17 − 8 = 9
neutrons. There are 8 electrons.
(d) Carbon-14 can also be represented as 14C. The atomic
number of carbon is 6, so there are 14 − 6 = 8 neutrons.
The number of electrons is 6.