Ch. 1: Atoms - faculty.sdmiramar.edufaculty.sdmiramar.edu/nsinkaset/powerpoints/Chapter01.pdf ·...
Transcript of Ch. 1: Atoms - faculty.sdmiramar.edufaculty.sdmiramar.edu/nsinkaset/powerpoints/Chapter01.pdf ·...
Ch. 1: Atoms
Dr. Namphol Sinkaset
Chem 200: General Chemistry I
I. Chapter Outline
I. Introduction
II. Particulate View of the World
III. The Scientific Approach
IV. History of the Atom
V. Subatomic Particles
VI. Atomic Mass
VII. Atoms and the Mole
I. Real-Life Legos®
• Everything is comprised of small parts
connected into a complex whole.
• The structure of the whole determines
its properties.
II. Particles
• We will approach chemistry with two
key principles in mind:
1. Matter is particulate.
2. Structure of particles determines
properties of matter.
• Chemistry seeks to understand
properties of matter by studying the
structure of particles that compose it.
II. Matter
• Matter is anything that occupies space
and has mass.
• Everything around you is composed of
matter – desk, book, air.
• Remember: matter is particulate.
II. Atoms and Molecules
• Atoms are the basic particles that
compose ordinary matter.
• Atoms can bind to one another in
specific arrangements to yield
molecules.
• For example, a water molecule is
comprised of 1 oxygen atom and 2
hydrogen atoms.
II. Structure and Properties
• Boils at 30 °C
• Feels like gasoline
• Doesn’t dissolve salt
• Boils at 100 °C
• Feels like water
• Dissolves salt
II. Classifying Matter
• Any sample of matter is called a
substance.
• Matter can be classified by state or by
composition.
• State determined by relative positions
and interactions of particles.
• Composition determined by types of
particles.
II. States of Matter
II. States of Matter
• solid: strong particle attractions, pack in fixed
locations, only vibrate in place, not
compressible
• liquid: slightly weaker particle attractions,
pack in non-fixed locations, fixed volume,
assume shape of container
• gas: weak particle attractions, free to move,
large distances between particles,
compressible
II. Composition of Matter
• Can also classify matter by the kinds of
particles out of which it is comprised.
• If there is only one type of particle, then
it is a pure substance.
• If there is more than one type of
particle, then it is a mixture.
II. Types of Matter
III. The Scientific Approach
• a.k.a. scientific method, is a flexible process
of creative thinking and testing aimed at an
objective
III. Differences Between
Hypothesis and Theory
• Hypothesis not thoroughly tested
• Theory more “developed”
• Hypothesis does not predict
• Experiments on hypothesis test hypothesis itself
• Experiments on theory test predictions of theory
• Theory can be expanded to many related situations
III. Differences Between
Hypothesis and Theory
• Compare the two statements below.
• “Methane reacts w/ oxygen to form
carbon dioxide and water.”
• “Hydrocarbons undergo a combustion
reaction w/ oxygen to form carbon
dioxide and water.”
IV. History of the Atom
• The Greeks were the first to wonder
about matter.
• Greek philosophers around 430 B.C.E.
debated what made up the world
around them.
• Leucippus and Democritus vs. Plato
and Aristotle
IV. Atomos vs. Fire, Air, Earth,
Water
IV. Revival of the Atom
• The idea of the atom was discarded and
forgotten about for almost 2000 years.
• In the late 18th and early 19th centuries,
three natural laws baffled everyone.
• John Dalton resurrected the idea of the
atom to explain what was observed.
IV. Law of Mass Conservation
• In a reaction, matter
is neither created
nor destroyed.
• Credit Antoine
Lavoisier (1789).
IV. Law of Mass Conservation
IV. Law of Definite Proportions
• All samples of a given compound have
the same proportions of constituent
elements.
Credit Joseph Proust (1797)
• e.g. Ammonia has 14.0 g N for every
3.0 g of H:
IV. Law of Multiple Proportions
• In 1804, John Dalton found that when two elements (A and B) form two different compounds, the masses of element B that combine with 1 g of element A can be expressed as a ratio of small whole numbers.
IV. Dalton’s Atomic Theory
• John Dalton revived the idea of the
atom to explain the natural laws that
had everyone perplexed.
• His atomic theory (1808) worked so well
that it was quickly accepted.
IV. Postulates of Dalton’s Theory
1. Each element is composed of tiny, indestructible particles called atoms.
2. All atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements.
3. Atoms combine in simple, whole-number ratios to form compounds.
4. Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms only change the way that they are bound together with other atoms.
IV. The Nuclear Atom
• Dalton’s theory treated atoms as
permanent, indestructible building
blocks that composed everything.
• A series of experiments were conducted
that led to a new view of the atom.
IV. Cathode Rays
• What conclusions about cathode rays can be
made from these experiments?
IV. Cathode Rays
• Using EM fields (late 1800s), J.J. Thomson measured the cathode ray particle’s mass to charge ratio.
• He estimated that cathode ray particles were about 2000 times lighter than a hydrogen atom.
• Result implies that atoms can be divided into smaller particles.
IV. Cathode Rays
• Using his famous oil drop experiment,
Robert Millikan (1909) calculated the
charge of a cathode ray particle.
• His value is w/in 1% of today’s accepted
value: -1.602 x 10-19 C.
• Mass was determined to be
9.109 x 10-28 g.
• Of course, cathode ray particles are
now known as electrons.
IV. Plum Pudding
• If electrons are in all matter, there must be positively-charged species as well.
• J.J. Thomson proposed the plum pudding model of the atom. Electron “raisins”
“Pudding” of positive charge
IV. The Role of Radioactivity
• Henri Becquerel and Marie Curie
discovered radioactivity by accident.
• Ernest Rutherford used radium, an
alpha (a) particle emitter.
• These a-particles are dense and have a
positive charge.
IV. Rutherford’s a-Particle
Experiment (1909)
IV. Conclusions from
Rutherford’s Experiment
• Most of an atom’s mass and all of its
positive charge exists in a nucleus.
• Most of an atom is empty space,
throughout which electrons are dispersed.
• By having equal numbers of protons and
electrons, an atom remains electrically
neutral.
• Note: neutrons discovered 20 years later.
IV. Rise of the Nuclear Atom
V. Subatomic Particles
• Therefore, all atoms are made up of
protons, neutrons, and electrons.
V. Atomic Number
• The atomic number (Z) of an element equals the # of protons in the nucleus
All atoms of an element have same, unique atomic number!!
• Protons are responsible for an atom’s identity.
• e.g. All carbon atoms have 6 protons and all uranium atoms have 92 protons.
V. Chemical Symbols
• Each element has a unique symbol.
• The symbol is either a 1 or 2
abbreviation of its name.
• e.g. carbon C; nitrogen N; chlorine
Cl; sodium Na; gold Au
V. Mass Number
• The mass number (A) is the total
number of protons and neutrons in the
nucleus.
• e.g. A carbon atom with 6 neutrons has
a mass number of 12.
V. Isotopes
• The # of protons determines the identity of the atom, but the # of neutrons has no effect.
• Thus, atoms of the same element can have different mass numbers.
• Since chemical properties are mainly due to e-, isotopes are almost identical chemically.
• Different isotopes of an element exist in certain percentages – natural abundances.
V. Depicting an Isotope
V. Sample Problem 1.1
1. What are the atomic number, mass
number, and symbol for the carbon
isotope with 7 neutrons?
2. How many protons and neutrons are
present in an atom of potassium-39?
V. Ions
• Atoms can lose or gain electrons and
become ions.
• Ions are charged particles.
• Positively-charged particles = cations.
e.g. Li Li+ + 1e-
• Negatively-charged particles = anions.
e.g. F + 1e- F-
VI. Atomic Mass
• Postulate #2 of Dalton’s atomic theory
stated that all atoms have the same
mass.
• With the existence of isotopes, this can’t
be true, but we can calculate an
average mass.
• The average mass of an element is
called the atomic mass.
VI. Atomic Mass
• Atomic masses are calculated using a
weighted average of all isotopes of an
element.
• The natural abundance is used to weight
each isotope in the calculation.
VI. Sample Problem 1.2
• Lithium has an atomic mass of 6.941
amu and has two naturally-occuring
isotopes: 6Li and 7Li. These isotopes
have isotopic masses of 6.0151 amu
and 7.0160 amu, respectively. What
are the percent natural abundances (to
2 decimal places) of these isotopes of
lithium?
VII. How Much vs. How Many
• Countable vs. not countable
• In the lab, we say, “How much water do
we need?”
• However, matter is particulate, so it is
countable.
• When matter interacts, it does so on a
particle by particle basis.
VII. Using Mass to “Count”
• We can relate counts to mass.
e.g. 21-30 shrimp vs. 8-10 shrimp
• In chemistry, we use the mole.
A mole is the amount of material that
contains 6.02214 × 1023 particles.
Defined by the # of atoms in exactly 12 g of
pure carbon-12.
This number is Avogadro’s number.
VII. Avogadro’s Number
• We can use Avogadro’s number as a conversion factor to calculate # of atoms.
6.022 x 1023 atoms
1 mole atoms 6.022 x 1023 atoms
1 mole atomsOR
VII. Grams and Moles
• The mass of 1 mole of atoms of an element is its molar mass. The value of an element’s molar mass in g/mole is numerically equal to the element’s atomic mass in amu (from periodic table).
• In general, one mole scales up to a “touchable” amount of an element or molecule.
VII. Mole of Atoms & Compounds
• 1 Fe atom weighs 55.85 amu, so 1 mole
of Fe atoms weighs 55.85 g.
• 1 O atom weighs 16.00 amu, so 1 mole
of O atoms weighs 16.00 g.
• Same applies for compounds.
1 molecule H2O weighs 18.02 amu, so 1
mole H2O weighs 18.02 g.
VII. Sample Problem 1.3
• If a silicon chip has a mass of 5.89 mg,
how many Si atoms are in the chip?