Download Atoms, molecules and ions

Atoms molecules and ions
Chapter 2
Conservation of mass and the Law
of definite proportions
• Democritus (460-370 BC) – proposed that elements are
composed of tiny particles , atoms from Greek atomos
meaning indivisible.
• Robert Boyle (1627 – 1691)- first to study chemistry as a
separate discipline and carry out rigorous chemical
experiments.
• Joseph Priestly (1733)- isolated oxygen and Antoine
Lavoiser(1743 – 1794)- oxygen key substance in
combustion, matter is neither created nor destroyed in a
chemical reaction.
• Joseph Proust (1754 – 1826)- different samples of a
pure chemical substance always contain the same
proportion of elements by mass, elements combine in
specific proportions, not just random proportions.
Dalton’s atomic theory and the law
of multiple proportions
•
1.
2.
3.
4.
John Dalton (1766 – 1844)- new theory of matter.
Elements are made up of tiny particles called atoms.
Each element is characterized by the mass of its
atoms. Atoms of the same elements have the same
mass, but atoms of different elements have different
masses.
Chemical combination of elements to make different
substances occurs when atoms join together in small
whole-number rations. (not fractions)
Chemical reactions only rearrange the way atoms are
combined, the atoms themselves don’t change.
Law of multiple proportions
• A theory must not only explain
observations but predict not yet know.
Dalton’s predicts LoMP.
• Elements can combine in different ways to
form different substances, whose mass
rations are small whole- number multiples
of each other.
Practice
• Methane and propane are both
constituents of natural gas. A sample of
methane contains 5.70 g of carbon atoms
and 1.90 g of hydrogen atoms combined in
a certain way, whereas a sample of
propane contains 4.47 g of carbon atoms
and 0.993 g of hydrogen atoms combined
in a different way. Show that the two
substances obey the law of multiple
proportions.
The structure of the atom: electrons
• J.J. Thomson (1856 – 1940) experimented
with cathode – ray tubes (sealed vacuum
tube with electrode covered with ZnS in it).
Cathode ray deflected by a magnet or
electrically charged plate. Beam is
produced at negative electrode and is
deflected to positive plate, must be tiny,
negative particles (electrons)
3 factors for electric field
• 1. the strength of the deflecting magnet or
electric field. (stronger – greater
deflection)
• 2. The size of the negative charge on the
electron. (larger charge – greater
deflection)
• 3. The mass of the electron. (lighter
particle – greater deflection)
Charge to mass ration
• e/m = 1.758820 X 108 C/g
e is the magnitude of the charge in coulomb C,
and m is mass in grams.
R. A. Millikan (1868 – 1953) devised a method
for measuring the mass of the electron. Oil
mist spayed into a chamber allowed to drop
between two plates. Mass from falling rate.
X-rays to oil made negative and voltage to
plates – oil suspended. e was calculated and
put above. 9.109 382 x 10 -28 g
The structure of atoms: protons
and neutrons
• Ernest Rutherford (1871 – 1937) directed
alpha particles (radiation – positive 2x’s
electron) at a thin gold foil. He found most
passed through (lost of space). Few were
deflected (even backwards). The mass in
the tiny nucleus. Atom’s diameter 10-10m
and nucleus 10 -15m (pea in center of
domed stadium) Protons mass – 1.672622
x 10-24 g, Neutron mass – 1.674927 x 10-24
Practice calculations using
atomic size
• Ordinary “lead” pencils actually are made
of a form of carbon called graphite. If a
pencil line is 0.35 mm wide and a diameter
of a carbon atom is 1.5 x 10-10 m, how
many atoms wide is the line?
Atomic number (z)
• Elements differ according to the number of
protons in their atoms’ nuclei.
• Number of protons = number of electrons
around atom’s nucleus.
• Mass number (A) = z + N (neutrons)
• Except for Hydrogen, most atoms contain
as many neutrons as protons
• Isotope – atoms with same number of
protons, but different number of neutron.
Isotope notation
• 12
• 6C
Top number is the mass number,
bottom number the atomic number.
Say carbon-12
Isotope practice
• Practice – the isotope of uranium used to
generate nuclear power is 235 U. How
many protons, neutrons, electrons.
• Element x is toxic to humans in high
concentration but is essential to life at low
concentrations. Identify element X, whose
atoms contain 24 protons, and write the
symbol for the isotope with 28 neutrons.
Atomic mass
• Atomic mass unit (amu) aka Dalton (Da) in
biological work. 1/12 the mass of C 12
and is equal to 1.600539 x 10-24 g. Protons
and neutrons are 1 amu.
• Weighted Average of all the isotopic
masses of the element’s naturally
occurring isotopes.
• C12 (98.89 % ) and C13 (1.11%). Carbon
mass 12.011
Calculating an atomic mass
• Chlorine has two naturally occuring
isotope: 35/17 Cl with a natural abundance
of 75.77% and an isotopic mass of 34.969
amu and 37/17 with a natural abundance
of 24.33% with a mass of 36.966 amu.
What is the atomic mass of chlorine?
Compounds
Chemical compound – when 2 or more different
elements combine in a specific way to create
a new material with different properties than
elements alone. Na (soft, silver) Cl ( green
gas) NaCl – table salt. Done by chemical
reaction. Formula – list symbol each element
in there. Subscript – tells the number of each
element.
Mixtures
• Simply blends of two or more substances
added together in arbitrary proportions,
without chemically changing the individual
substances. Be separated by physical
means.
• Heterogenous – not uniform mixture,
salad, muddy water, separates when still
• Homogenous – uniform, constant
composition. Air, salt water, koolaid
Covalent bonds: molecules
• 2 elements share (usually 2) electrons.
Like a tug-of-war.
• Molecules form this way like: HCl, H2O,
NH3 (ammonia), CO2.
• Model using ball and stick (specific
covalent bond) or space filling (overall
molecular shape)
• Structural formula show specific
connection between atoms. H-H
formulas
• Structural formulas useful in organic
chemistry.
• Some elements (right corner of periodic
table) exists as molecule (H2, N2, O2, F2,
Cl2, Br2, I2) are diatomic molecules.
• Practice: Propane has a structure in which
the 3 C atoms are bonded and each end C
is bonded to 3 H and center C is to 2 H.
Draw the structural formula.
Ionic bonds
• Transfer of 1 or more electrons to another
usually between metal and nonmetal.
• When atoms lose an electron, it becomes
more positive by 1, called cations, metals
• When atoms gain an electron, it becomes
more negative by 1, called anions,
nonmetals
• Practice: Show the formation of NaCl.
ions
• Can’t really talk about discrete Na+Clmolecules, but rather an ionic solid. Table
salt crystal – cube shape
• Covalent bonded groups of atoms called
polyatomic ions have a charge form ionic
bonds. When more than one are in a
formula, put () around.
• Practice: which is ionic, which molecular
a) BaF2 2) SF4 3) PH3 4) CH3OH
Acids and bases
• Acid - Hydrogen (H+) cation in water (aq)
and the anion nonmental / polyatomic ion;
hydrochloric (HCl), Nitric (HNO3), sulfuric
(H2SO4), Phosphoric acid (H3PO4); the a
• Base - Hydroxide anion (OH-) in water and
the metal cation; sodium hydroxide (NaOH
or lye or caustic soda), Potassium
hydroxide (KOH caustic potash) and
barium hydroxide [Ba (OH)2]
Acid / base practice
• Which of the following compounds are
acids, and which are bases, explain.
• A) HF
• B) Ca(OH)2
• C) LiOH
• D) HCN
Naming Binary Ionic
compounds
• Name must defines it uniquely but also
allow chemists (and computers) to know
the chemical structure. Organics more
complicated. Use simple for now.
• Postive ion takes the name of the element
(metal), negative side (nonmetal) gets the
–ide ending.
• Ions formed by main groups: 1 (+1), 2 (+2)
3 (+3), 4 (metals) (+4 or +2), 5 (-3), 6 (-2),
7 (-1) Group 8 = 0
Transitional naming
• Transitional carry multiple charges. Fe+2,
Fe+3. distinguish by using roman numerals
in (). Iron (II), Iron (III). Older forms use
ferrous (lower charge), ferric (higher
charge).
• Neutral compound: total number of
positive charge = total number of negative.
One can figure out number of +, by
counting the number of negatives on
anions.
Practice
• Give systematic names for the following:
a) BaCl2 b) CrCl3 c) PbS d) Fe2O3
Write the formulas for the following:
a) Magnesium fluoride b) Tin (IV) oxide
c) Iron (III) sulfide
Naming binary molecular
compounds
• Assume one to be more cation (farther
left) like – 1st word and just name, 2nd to
be more anion like. Tell how many there
are by the subscript using prefixes. Mono
– 1, di – 2, tri – 3, tetra – 4, penta – 5,
hexa – 6, hepta – 7, octa – 8, nona – 9,
dec – 10.
• If prefix ends in a or o and anion name
begins with a vowel, drop a or o to avoid 2
vowels together.
Practice
Give systematic names for the following:
• a) PCl3
• b) N2O3
• c) P4O7
• d) BrF3
Write formulas for:
a) Carbon dioxide
b) Tetraphoshorus hexachloride
Naming compounds with
polyatomic ions
• Named same way as binary ionic.
Ba(NO3)2 is called barium nitrate.
• Most have ite or ate ending. (3 , –ide),
several form oxoanions (with oxygen). If 2
forms then more O2 gets -ate, less –ite.
More than 2 O2 hypo – for least, then –ite,
-ate, per – for most
• Will be given a list to memorize.
Practice with polyatomic
• Give systematic names for the following:
a) LiNO3
b) KHSO4
c) CuCO3
d) Fe(ClO4)3
Write the formulas
a) Potassium hypochlorite
b) Silver (I) chromate
c) Iron (III) carbonate
Naming acids
• Most are oxoacid – contain oxygen in
addition to hydrogen. –ite gets ous acid
and –ate gets ic acid.
• Binary acid use hydro anion ic acid. HCl –
hydrochloric acid
• Practice: name the following acids.
a)HBrO (aq)
b)HCN (aq)