Download Chapter 2 Notes - Waterford Public Schools

WHAT IS MATTER?
A BRIEF REVIEW OF THE CLASSIFICATION, PROPERTIES,
AND SEPARATION OF MATTER
MIND CATALYST
• For each image on page 25:
• Categorize the substance as either an element, a
compound, or a mixture.
• Briefly explain why you categorized it the way you did.
• Draw a particle diagram that describes your
categorization. What does an element look like on
the molecular level? What about a compound or a
mixture? Use labels if necessary.
• If the substance is undergoing a process, explain
whether it is a physical or chemical process. Justify
your answer.
MIND CATALYST
ATOMIC STRUCTURE AND
THE PERIODIC TABLE
MODERN VIEW OF ATOMIC STRUCTURE AND PERIODIC
TABLE TRENDS
ATOMS AND ATOMIC THEORY
• There are five postulates that describe the atomic
theory of matter:
1.
2.
3.
4.
5.
Matter consists of indivisible atoms.
All of the atoms of a given chemical element are identical
in mass and in all other properties.
Different chemical elements have different kinds of atoms,
and such atoms have different masses.
In an ordinary chemical reaction, atoms move from one
substance to another, but no atom of element disappears
or is changed into an atom of another element.
The formation of a compound from its elements occurs
through the combination of atoms of unlike elements in
small whole number ratios.
FUNDAMENTAL LAWS OF MATTER
• There are also three fundamental laws of matter:
• Law of conservation of mass
• Matter is conserved in chemical reactions
• Law of constant composition
• Pure water has the same composition everywhere
• Law of multiple proportions
• Compare Cr2O3 to CrO3
• The ratio of Cr:O between the two compounds is a small whole
number
THE FORMAL DEFINITIONS OF THE
THREE LAWS
• Law of conservation of mass
• In every chemical reaction equal quantity of matter exists
before and after the reaction.
• Verified by Antoine Lavoisier
• Law of constant composition
• In a given compound, the proportions by mass of the elements
that compose it are fixed, independent of origin of the
compound or the type of the preparation.
• Discovered by Joseph Proust
• Law of multiple proportions
• When two elements form a series of compounds, the ratios of
the masses of the second element that combine with 1 gram
of the first element can always be reduced to small whole
numbers.
• Discovered by John Dalton
BUT THERE WAS A PROBLEM…
• Atomic theory raised more questions than it
answered. For example:
• Could atoms be broken down into smaller particles?
• Over a 100 years after the atomic theory of
proposed, the answers were provided by many
experiments…
ATOMIC THEORY DEVELOPMENT
• Democritus (400 B.C)
• First used the term “atom” - atomos
• John Dalton (1803)
• Credited with being the “Father of Modern Chemistry”
• Developed 4 postulates
• Elements are made up of tiny particles called atoms
• Atoms of a given element are identical; atoms of different
elements are different
• Chemical compounds are formed when atoms of different
elements combine with each other
• Chemical reactions involve reorganization of atoms; the atoms
themselves are not changed in a chemical reaction
ATOMIC THEORY DEVELOPMENT
• J.J Thomson (1897)
• Credited with the discovery of the electron
• Max Planck (1900)
• Described packets of energy called quanta
• Albert Einstein (1905)
• Described photoelectric effect
• Described wave-particle duality of radiation
• Robert Millikan (1909)
• Discovers magnitude of electron charge and mass of the
electron through oil-drop experiments
• Ernest Rutherford (1911)
• Gold-foil experiment
• Proposed the nuclear atom
• Atom is mostly empty space
ATOMIC THEORY DEVELOPMENT
• Neils Bohr (1913)
• Electrons move in fixed orbits
• Model only worked for hydrogen atoms
• Louis de Broglie (1923)
• Proposes particle wave behavior of electron
• Particle-wave duality
• Erwin Schrodinger (1926)
• Formulates an equation to determine probability of electron
location
• Quantum Theory
• James Chadwick (1932)
• Discovers neutron
SO, WHAT ARE THE
COMPONENTS OF ATOMS?
A SUMMARY OF ATOMIC STRUCTURE
• The atom consists of positive, negative, and neutral entities called
protons, electrons, and neutrons
• Protons determine element’s identity
• # of protons is unique for each element
• Electrons determine element’s chemical properties
• Neutrons act as a “glue” for the protons to minimize charge repulsions
• Protons and neutrons are located in the nucleus of the atom, which is
small
• Therefore, the nucleus is positively-charged
• Electrons are located outside of the nucleus at an average distance
of 10-8 cm
• Therefore, the electron cloud is negatively-charged
NUCLEUS SIZE VS. ELECTRON CLOUD
SIZE
• Most of the
volume of the
atom is due to
electrons
MASSES OF THE PROTON, NEUTRON,
ELECTRON
• Most of the mass of the atom is due to the
protons and neutrons within the nucleus
IMPORTANT DEFINITIONS TO KNOW
• Atomic number (Z)
• Number of protons in the nucleus
• Mass number (A)
• Total number of nucleons in the nucleus (i.e., protons and
neutrons)
• Isotopes have the same Z but different A!
ISOTOPES OF CARBON
PRACTICE!
Nuclear
symbol
31P
Charge
Number of
protons
Number of
neutrons
0
9
10
0
+3
Number of
electrons
16
27
30
16
16
18
ATOMIC MASS VS. MASS NUMBER
• Look on your periodic table at the mass of carbon
• The red number is NOT the mass number!!
• It is the average atomic mass of ALL isotopes of carbon that
are known to exist
• So, the element carbon on the periodic table is actually a
mixture of the isotopes carbon-12, carbon-13, and carbon14
• Therefore, the atomic mass on the periodic table is
a weighted average of the atomic mass of all
isotopes of that particular element
•
One atomic mass unit is equal to 1.66 x 10-24 grams, which is
1/12th the mass of a 12C atom
DETERMINING ATOMIC MASSES
• Atomic masses can be determined to highly precise values
by using a mass spectrometer
• The mass spectrometer separates matter based on its mass
and charge
• Atoms are ionized at low pressure in the gas phase
• The cations that form are accelerated toward a magnetic field
• The extent to which the cation beam is deflected is inversely
related to the mass of the cation
• The resulting data is plotted with abundance on the y-axis and
mass on the x-axis
• The mass spectrometer is used to determine:
• The types of isotopes present in an element
• The exact atomic masses of these isotopes
• The relative amount of each isotope present
CALCULATING AVERAGE ATOMIC
MASS FOR AN ELEMENT
• One can calculate the average atomic weight of
an element if the abundance of each isotope for
that element is known
Average Atomic Mass
% natural abundance
=
∙ atomic mass1
100
1
% natural abundance
+
∙ atomic mass2 …
100
2
EXAMPLE
•The average atomic mass of naturally occurring neon is 20.18 amu. There
are two common isotopes of naturally occurring neon as indicated in the
table below.
Isotope
Ne-20
Ne-22
(i)
Mass (amu)
19.99
21.99
Using the information above, calculate the percent abundance of
each isotope.
(ii)
Calculate the number of Ne-22 atoms in a 12.55 g sample of
naturally occurring neon.
ATOMS, ELEMENTS, & THE
PERIODIC TABLE
WHY IS THE PERIODIC TABLE
IMPORTANT?
• The Periodic Table is used to organize the 114
elements in a meaningful way
• Arranged by increasing atomic number
A BRIEF REVIEW OF THE ORGANIZATION OF
THE PERIODIC TABLE
• Columns in the periodic table are called groups
• Numbered from 1A to 8A or 1 to 18
• Atoms with similar properties appear in groups or families
• They are similar because they all have the same number of
valence (outer shell) electrons, which governs their chemical
behavior
• Remember, valence electrons are electrons in the highest-numbered sand p- orbitals!
• Rows in the periodic table are called periods
• Elements of the same period have the same
number of energy levels
• As you move across a period, the number of electrons and protons
increases, leading to increase in atomic number
• Elements within the same period do not generally show similarity in
properties, except d-block and f-block (lanthanides) elements
THE PERIODIC TABLE
FAMILIES OF THE PERIODIC TABLE
• Some of the groups in the periodic table are
given special names, and are called families
• These names indicate the similarities of
chemical properties between group members
as a result of same number of valence
electrons:
•
•
•
•
•
Group 1:Alkali metals
Group 2: Alkaline earth metals
Group 3-12: Transition metals
Group 17: Halogens
Group 18: Noble gases
A BRIEF REVIEW OF THE
ORGANIZATION OF THE PERIODIC
TABLE
• Metals are located on the left hand side of the Periodic Table
• Most of the elements are metals
• The metals include all of the transition metals as well as post transition
metals
• Metals are solids (except mercury), conduct electricity, are ductile, are
malleable and can form alloys
• Non-metals are located in the top right hand side of the Periodic Table
• They have wide variety of properties:
• Some are solids, bromine is a liquid; and some are gases at room temperature.
With the exception of graphite, a form of carbon, nonmetals do not conduct
electricity.
• Elements with properties similar to both metals and non-metals are called
metalloids and are located at the interface between the metals and nonmetals
•
•
•
•
•
•
B
Si
Ge
As
Sb
Te
USING THE PERIODIC TABLE TO
PREDICT ION CHARGE
• The number of electrons an atom loses or gains is
related to its position on the periodic table
• Metals tend to form cations whereas non-metals
tend to form anions
WHAT IS A CATION?
• When an atom or molecule loses electrons, it
becomes positively charged
• For example, when Na loses an electron it
becomes Na+
• Positively charged ions are called cations
WHAT IS AN ANION?
• When an atom or molecule gains electrons, it
becomes negatively charged
• For example when Cl gains an electron it
becomes Cl• Negatively charged ions are called anions
• An atom / molecule can lose or gain more than 1
electron!
UNDERSTANDING ELECTRONS &
CHEMICAL REACTIVITY
MOLECULAR & IONIC COMPOUNDS
• In an ordinary chemical reaction, the nucleus of
each atom remains unchanged.
• However, electrons can be:
• Added to atoms by transfer from other atoms,
• Lost by transfer to other atoms
• Or shared with other atoms.
FORMATION OF IONIC COMPOUNDS
• Occurs when a metal loses all of its valence
electrons to a non-metal
• Metal becomes positively-charged (cation)
• Non-metal becomes negatively-charged (anion)
• Opposite charges hold the compound together!
• Electrostatic attraction
• When ions are combined together to form IONIC
compounds, the overall charge of the compound
that results MUST BE ZERO (neutral)!
WHAT ARE COVALENT COMPOUNDS?
• Remember, ionic
compounds generally
involve a metal cation (or
ammonium) and a nonmetal anion (or
polyatomic anion) being
held together by an
electrostatic attraction
• Compounds are called
formula units
• Covalent compounds
consist of two nonmetals
sharing electrons
• Compounds are called
molecules
ACIDS AND THEIR FORMULAS
• Acids are an important class of hydrogen-containing
compounds, and they are named in a special way
• A simple definition of an acid is a substance which
produces H+ ions in water
• Note – In order for a substance to be an acid instead of a gas (i.e.
HCl), binary acids must be aqueous or dissolved in water
• Ternary (or oxy-) acids are ALWAYS aqueous
• When we encounter the chemical formula for an acid, it
will be written with an H as the first element
• Examples:
• HCl (aq) – hydrochloric acid
• HNO3 (aq)– nitric acid
• H2SO4 (aq) – sulfuric acid
CHEMICAL NOMENCLATURE
• Know how to properly name and write the formula
for:
• Binary ionic compounds (Type I)
• Contains a metal and non-metal
• Metal only forms a single type of cation
• Example – NaCl
• Binary ionic compounds (Type II)
• Contains a metal and non-metal
• Contains a metal that forms more than one type of positive ion
so charge must be specified using Roman numerals
• Example – HgO
• Ionic compounds with polyatomic ions
• Must memorize list of common polyatomic ions
CHEMICAL NOMENCLATURE
• Know how to properly name and write the formula
for:
• Binary covalent compounds (Type III)
• Formed between two non-metals
• Use prefixes!
• Example – P4O10
• Acids
• When dissolved in water, some molecules produce a solution
containing free H+ ions
• Example – HClO4
DIFFERENT FORMULA
REPRESENTATIONS
DIFFERENT FORMULA
REPRESENTATIONS
• Molecules can be represented several different ways.
•
Molecular formula in which the symbols for the elements are used to
indicate the types of atoms present and subscripts are used to indicate the
relative number of atoms.
•
Empirical formula shows the types of atoms present and the simplest whole
number ratio of the number of atoms in the compound.
•
Structural formula of a substance shows which attached to which within
the molecule.
•
Condensed structural formula suggests the bonding pattern in the
molecule and highlights the presence of a reactive group of atoms within
the molecule.
EMPIRICAL FORMULA
• Benzene, C6H6, is produced during oil refining and has many
industrial uses. A benzene molecule can be represented as (a)
a structural formula, (b) a ball-and-stick model, and (c) a
space-filling model. (d) Benzene is a clear liquid. (credit d:
modification of work by Sahar Atwa)
• (a) Vinegar contains acetic acid, C2H4O2, which
has an empirical formula of CH2O. It can be
represented as (b) a structural formula and (c) as
a ball-and-stick model. (credit a: modification of
work by “HomeSpot HQ”/Flickr)
• Molecules of (a) acetic acid and methyl formate
(b) are structural isomers; they have the same
formula (C2H4O2) but different structures (and
therefore different chemical properties).