Overview of the
Basics
CHAPTER 1-3 Review
Chemistry: The Molecular Nature of Matter, 6th edition
By Jesperson, Brady, & Hyslop
CHAPTER 1-3 Review
Learning Objectives
 Scientific Method
 Matter: definition, elements, compounds, mixtures, changes/properties
 Atomic Theory
 Law of definite proportions
 Law of conservation of mass
 Chemical formulas
 Chemical equations
 Balancing
 Measurements: units, conversions, uncertainty
 Significant Figures
 Density
 Subatomic particles
 Atomic #, mass #, atomic weights
 Periodic Table
 Ionic Compounds, hydrates, molecular compounds
 Basic nomenclature
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
1
Scientific Method
1. Make observations/collect data
•
•
•
•
2.
Law or Scientific Law
•
•
•
•
3.
Usually an equation
Based on results of many experiments
Only states what happens
Does not explain why they happen
Hypothesis
•
•
•
•
•
4.
Empirical fact
Something we see, hear, taste, feel, or smell
Something we can measure
Organize data so we can see relationships
Mental picture that explains observed laws
Tentative explanation of data
Make predictions
Leads to further tests
Go to laboratory and perform experiments
Theory
•
•
•
•
Tested explanation of how nature behaves
Devise further tests
Depending on results, may have to modify
theory
Can never prove theory is absolutely correct
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
1
Elements
• Substances that can’t be decomposed into
simpler materials by chemical reactions
• Substances composed of only one type of atom
• Simplest forms of matter that we can work with
directly
• More complex substances composed of elements
in various combinations
diamond = carbon
gold
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
sulfur
4
Chapter
1
Elements
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
1
http://ridenourmhs.wikispaces.com/ESUnit2
Classification of Matter
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
2
Chemical vs Physical Properties
Physical properties
Can be observed without changing
chemical makeup of substance
Solids:
Fixed shape and volume
Particles are close together
Liquids:
Fixed volume, but take container shape
Particles are close together
Gases:
Expand to fill entire container
Particles separated by lots of space
Chemical properties
• Chemical change or reaction that
substance undergoes
• Chemicals interact to form entirely
different
substances with different chemical
and physical properties
• Describe behavior of matter that
leads to formation of new substance
• “Reactivity" of substance
e.g. Iron rusting
– Iron interacts with oxygen to
form a new substance.
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
1
Atomic Theory
Developed by John Dalton to explain Law of
Conservation of Mass & Law of Definite Proportions
1. Matter consists of tiny particles called atoms.
2. Atoms are indestructible.
•
In chemical reactions, atoms rearrange but do
not break apart.
3. In any sample of a pure element, all atoms are
identical in mass and other properties.
4. Atoms of different elements differ in mass and other
properties.
5. In a given compound, constituent atoms
are always present in same fixed numerical ratio.
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
1
Law of Definite Proportions
Atoms react as Whole particles.
When two elements form more than one compound,
different masses of one element that combine with
same mass of other element are always in ratio of
small whole numbers.
e.g. Fool’s gold, pyrite, iron(III) sulfide
Mass ratio always
1.00 g of iron to 0.574 g of sulfur
e.g. Water
Mass ratio always: 8 g O to 1 g H
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
1
Law of Conservation of Mass
sulfur sulfur
dioxide trioxide
Mass S
Mass O
32.06 g
32.06 g
32.00 g
48.00 g
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
1
Molecules and Chemical Formulas
Atoms combine to form more complex substances = Molecules
Chemical Formulas:
• Specify composition of substance
• Chemical symbols represent atoms of elements present
• Subscripts:
– Given after chemical symbol
– Represents relative numbers of each type of atom
Example:
Fe2O3 : iron and oxygen in 2:3 ratio
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
1
Hydrates
• Crystals that contain water molecules
e.g. Plaster: CaSO4∙2H2O calcium sulfate dihydrate
– Water is not tightly held
• Dehydration
– Removal of water by heating
– Remaining solid is anhydrous (without water)
White = CuSO4
Blue =
CuSO4 •5H2O
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
1
Depicting Molecules
CH4
methane
H
H
C
H
H
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
1
Chemical Equations
• Use chemical symbols and formulas to represent
reactants and products.
– Reactants on left hand side
– Products on right hand side
– Arrow () means “reacts to yield”
e.g. CH4 + 2O2  CO2 + 2H2O
– Coefficients
• Numbers in front of formulas
• Indicate how many of each type of molecule
reacted or formed
– Equation reads “methane and oxygen react to
yield carbon dioxide and water”
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
1
Conservation of Mass in Reactions
• Mass can neither be created nor destroyed
• This means that there are the same number of each type
of atom in reactants and in products of reaction
– If number of atoms same, then mass also same
CH4 + 2O2

4H + 4O + C
=
CO2 + 2H2O
4H + 4O + C
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
1
Ex.
Balanced Chemical Equations
2C4H10 + 13O2  8CO2 + 10H2O
4 C and 10 H
per molecule
Ex.
2 O per
molecule
2 H and 1 O
1 C and 2
per
O per
molecule
molecule
2C4H10 + 13O2  8CO2 + 10H2O
2
molecules
of C4H10
13 molecules
of O2
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
8
molecules
of CO2
10 molecules
of C4H10
16
Chapter
2
Intensive vs Extensive Properties
Intensive properties
– Independent of sample size
– Used to identify substances
e.g. Color
Density
Boiling point
Melting point
Chemical reactivity
Extensive properties
– Depend on sample size
e.g. volume and mass
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
2
Measurements
1. Measurements involve comparison
– Always measure relative to reference
e.g. Foot, meter, kilogram
– Measurement = number + unit
e.g. Distance between 2 points = 25
• What unit? inches, feet, yards, miles
• Meaningless without units
2. Measurements are inexact
– Measuring involves estimation
– Always have uncertainty
– The observer and instrument have inherent
physical limitations
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
2
International System of Units
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
2
International System of Units
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
2
International System of Units
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
2
Decimal Multipliers
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
2
4 Common Lab Measurements
1. Distance (d )
Centimeter (cm)
1 cm = 10–2 m = 0.01 m
Millimeter (mm)
1 mm = 10–3 m = 0.001 m
2. Volume (V)
1 L = 1000 mL
1 mL = 1 cm3
3. Mass (m)
1 g = 0.1000 kg =
1
g
1000
4. Temperature (T)
273.15 K = 0°C
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
2
Uncertainty in Measurements
Measurements all inexact
Limitations of reading instrument
Example: Consider two Celsius thermometers
• Left thermometer has markings every 1˚C
– T between 24 °C and 25 °C
– About 3/10 of way between marks
– Can estimate to 0.1 °C = uncertainty
– T = 24.3  0.1 °C
• Right thermometer has markings every 0.1 °C
– T reading between 24.3 °C and 24.4 °C
– Can estimate 0.01 °C
– T = 24.32  0.01 °C
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
Significant Figures
2
Scientific convention: All digits in measurement up to and
including first estimated digit are significant.
1. All non-zero numbers are significant.
e.g. 3.456
has 4 sig. figs.
2. Zeros between non-zero numbers
are significant.
e.g. 20,089
or
2.0089 × 104
has 5 sig. figs
3. Trailing zeros always count as significant if number
has decimal point
e.g. 500.
or
5.00 × 102
has 3 sig. figs
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
2
Significant Figures
4. Final zeros on number without decimal point are
NOT significant
e.g. 104,956,000
or
1.04956 × 108
has 6 sig. figs.
5. Final zeros to right of decimal point are significant
e.g. 3.00 has 3 sig. figs.
6. Leading zeros, to left of first nonzero digit, are never
counted as significant
e.g. 0.00012 or 1.2 × 10–4
has 2 sig. figs.
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
2
Significant Figures: Rounding
1. If digit to be dropped is greater than 5, last
remaining digit is rounded up.
e.g. 3.677 is rounded up to 3.68
2. If number to be dropped is less than 5, last
remaining digit stays the same.
e.g. 6.632 is rounded to 6.63
3. If number to be dropped is exactly 5, then if digit to
left of 5 is
a. Even, it remains the same.
e.g. 6.65 is rounded to 6.6
b. Odd, it rounds up.
e.g. 6.35 is rounded to 6.4
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
Significant Figures: Calculations
2
Multiplication and Division
• Number of significant figures in answer = number of
significant figures in least precise measurement
e.g. 10.54 × 31.4 × 16.987
4 sig. figs. × 3 sig. figs. × 5 sig. figs. = 3 sig. figs.
Addition and Subtraction
12.9753
• Answer has same number
of decimal places as
319.5
quantity with fewest
+ 4.398
number of decimal
336.9
places.
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
4 decimal places
1 decimal place
3 decimal places
1 decimal place
28
Chapter
2
Significant Figures: Exact Numbers
• Numbers that come from definitions
– 12 in. = 1 ft
– 60 s = 1 min
• Numbers that come from direct count
– Number of people in small room
• Have no uncertainty
• Assume they have infinite number of significant figures.
• Do not affect number of significant figures in
multiplication or division
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
2
Scientific Notation
• Clearest way to present number of significant figures
unambiguously
– Report number between 1 and 10 followed by
correct power of 10
– Indicates only significant digits
e.g. 75,000 people attend a concert
– If a rough estimate
• Uncertainty 1000 people
• 7.5 × 104
– If number estimated from aerial photograph
• Uncertainty 100 people
• 7.50 × 104
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
2
Accuracy & Precision
Accuracy
– How close measurement is to true or
accepted true value
• Measuring device must be calibrated
with standard reference to give
correct value
Precision
– How well set of repeated
measurements of same quantity
agree with each other
– More significant figures equals
more precise measurement
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
2
Dimensional Analysis
• Also called the Factor Label Method
• Not all calculations use specific equation
• Use units (dimensions) to analyze problem
Conversion Factor
• Fraction formed from valid equality or
equivalence between units
• Used to switch from one system of
measurement and units to another
Given
× Conversion =
Quantity
Factor
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
Desired
Quantity
32
Chapter
2
Dimensional Analysis
Example: Convert 0.097 m to mm.
• Relationship is 1 mm = 1 × 10–3 m
• Can make two conversion factors
1  10 3 m
1 mm
1 mm
1  10  3 m
• Since going from m to mm use one on left.
0.097 m 
1 mm
1  10  3 m
= 97 cm
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
2
Density
• Ratio of object’s mass to its volume
mass
density 
volume
m
d=
V
• Intensive property (size independent)
– Determined by taking ratio of two extensive properties (size
dependent)
– Frequently ratio of two size dependent properties leads to size
independent property
• Density useful to transfer between mass and volume
of substance
• Density decreases slightly as temperature increases
• Units: g/mL or g/cm3
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
3
Discovery of electron
mass and charge
Millikan Oil Drop expt
Discovery of Subatomic Particles in the late
1800’s and early 1900s
Discovery of the Nucleus
Rutherford
Alpha scattering expt
Discovery of Protons
1918 Rutherford
Mass spectrometer
Discovery of the Electron
1897 Thomson
Cathode ray tube expt
Discovery of Neutron:
1932 Chadwick
Rutherford Nuclear Atom
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Properties of Subatomic Particles
Nucleus
(protons +
neutrons)
 Three kinds of
subatomic particles of
principal interest to
chemists
Particle
Mass (g)
Electrons
Electrical
Charge
Symbol
Electron 9.10939  10–28
–1
1.67264  10–24
+1
0
1 e
1
1
1 H, 1 p
0
1
0n
Proton
Neutron 1.67495  10–24
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
Atomic Notation
3
Atomic number (Z) = Number of protons that
atom has in nucleus
Isotopes = Atoms of same element with different masses
1
– Same number of protons (1 p )
– Different number of neutrons ( 10n )
Isotope Mass number (A)
– A = (number of protons)+(number of neutrons) = Z + N
– For charge neutrality, number of electrons and protons
must be equal
Atomic Symbols = Summarize information about subatomic
particles
– Every isotope defined by two numbers Z and A
Ex. What is the atomic symbol for helium?
4
He has 2 e–, 2 n and 2 p Z = 2, A = 4
2 He
37
A
Z
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Isotopes
• Most elements are mixtures of two or more stable isotopes
• Each isotope has slightly different mass
• Chemically, isotopes have virtually identical chemical
properties
• Relative proportions of different isotopes are essentially
constant
• Isotopes distinguished by mass number (A):
e.g.
– Three isotopes of hydrogen (H)
– Four isotopes of iron (Fe)
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Carbon-12 Atomic Mass Scale
• Need uniform mass scale for atoms
Atomic mass units (symbol u)
– Based on carbon:
• 1 atom of carbon-12 = 12 u (exactly)
• 1 u = 1/12 mass 1 atom of carbon-12 (exactly)
Why was 12C selected?
– Common
– Most abundant isotope of carbon
– All atomic masses of all other elements ~ whole
numbers
– Lightest element, H, has mass ~1 u
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Calculating Atomic Mass
• Generally, elements are mixtures of isotopes
e.g. Hydrogen
Isotope
Mass
% Abundance
1H
1.007825 u
99.985
2H
2.0140 u
0.015
How do we define atomic mass?
– Average of masses of all stable isotopes of given element
How do we calculate average atomic mass?
– Weighted average
– Use isotopic abundances and isotopic masses
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Periodic Table
– Summarizes periodic properties of elements
Early Versions of Periodic Tables
– Arranged by increasing atomic mass
– Mendeleev (Russian) and Meyer (German) in 1869
– Noted repeating (periodic) properties
Modern Periodic Table
– Arranged by increasing atomic number (Z ):
– Rows called periods
– Columns called groups or families
• Identified by numbers
• 1 – 18 standard international
• 1A – 8A longer columns and 1B – 8B shorter columns
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Periodic Table
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Periodic Table
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Periodic Table Groups
1A
2A
B
B
7A
8A
Alkali
Metals
Alkaline
Earth
Metals
Transition
Metals
Lanthanide
& Actinide
Halogens
Nobel
Gases
Very reactive Reactive
Metals
Metals
except for H
Form ions
with several
different
charges
(oxidation
states)
+1 ions
React with
Oxygen to
form
compounds
that dissolve
into alkaline
solutions in
water
+2 ions
Oxygen
compounds
are strongly
alkaline
Many are not
water soluble
Tend to form
+2 and +3
ions
Lanthanides
58 – 71
Actinides
90 – 103
Reactive
Inert
Form
diatomic
molecules in
elemental
state
Heavier
elements
have limited
reactivity
-1 ions
Actinides are Salts with
radioactive
alkali metals
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
Do not form
ions
Monoatomic
gases
44
Chapter
3
Metals, Nonmetals, and Metalloids
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
45
Chapter
3
Metals, Nonmetals, and Metalloids
Metals
• Metallic luster,
malleable, ductile,
hardness variable
• Conduct heat and
electricity
• Solids at room
temperature with the
exception of Hg
Nonmetals
• Brittle
• Insulators, nonconductors of
electricity and heat
• Chemical reactivity
varies
Metalloids
• Metallic shine but
brittle
• Semiconductors:
conduct electricity but
not as well as metals:
examples are silicon
and germanium
• Exist mostly as
compounds rather
then pure elements
• Chemical reactivity
• Many are gases,
varies greatly: Au, Pt
some are solids at
unreactive while Na, K
room temp, only Br2 is
very reactive
a liquid.
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Ions and Ionic Compounds
Ions
– Transfer of one or more electrons from one
atom to another
– Form electrically charged particles
Ionic compound
– Compound composed of ions
– Formed from metal and nonmetal
– Infinite array of alternating Na+ and Cl– ions
Formula unit
– Smallest neutral unit of ionic compound
– Smallest whole-number ratio of ions
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Ions and Ionic Compounds
Metal + Non-metal  ionic compound
2Na(s) + Cl2(g)  2NaCl(s)
Michael Watson
Na + Cl
Richard Megna/Fundamental Photographs
Richard Megna/Fundamental Photographs
Na+ + Cl
NaCl(s)
e
Anions = Negatively charged ions Cations = Positively charged ions
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Ions and Ionic Compounds
Electrical conductivity requires charge movement
Ionic compounds:
– Do not conduct electricity in solid state
– Do conduct electricity in liquid and aqueous states where ions are free
to move
Molecular compounds:
– Do not conduct electricity in any state
– Molecules are comprised of uncharged particles
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Ions and Ionic Compounds
Negative (–) charge on anion = number of spaces you have to move to
right to get to noble gas
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Rules for Writing Ionic Formulas
1. Cation given first in formula
2. Subscripts in formula must produce electrically neutral
formula unit
3. Subscripts must be smallest whole numbers possible
–
Divide by 2 if all subscripts are even
–
May have to repeat several times
4. Charges on ions not included in finished formula unit of
substance
–
If no subscript, then 1 implied
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Determining Ionic Formulas
“Criss-cross” rule
– Make magnitude of charge on one ion into subscript for
other
– When doing this, make sure that subscripts are reduced
to lowest whole number.
Ex. What is the formula of ionic compound formed between
aluminum and oxygen ions?
Al3+ O2–
Al2O3
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Transition Metal and Post-Transition Metal Ions
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
53
Chapter
3
Polyatomic
Ions
Example: What is the formula
of the ionic compound formed
between ammonium and
phosphate ions?
Ammonium = NH4+
Phosphate = PO43–
(NH4)+ (PO4)3–
(NH4)3PO4
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Nomenclature
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
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Chapter
3
Nomenclature: Ionic Compounds
Cations:
– Metal that forms only one positive ion
• Cation name = English name for metal
– Na+
sodium
– Ca2+ calcium
– Metal that forms more than one positive ion
– Use Stock System
• Cation name = English name followed by numerical
value of charge written as Roman numeral in
parentheses (no spaces)
• Transition metal
– Cr2+
chromium(II)
Cr3+
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
chromium(III)
56
Chapter
3
Nomenclature: Ionic Compounds
Anions:
– Monatomic anions named by adding
“–ide” suffix to stem name for element
– Polyatomic ions use names in Table 3.5
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
Nomenclature: Hydrates
3
• Ionic compounds
– Crystals contain water molecules
– Fixed proportions relative to ionic substance
• Naming
– Name ionic compound
– Give number of water molecules in formula
using Greek prefixes
monoditritetrapenta-
=
=
=
=
=
1
2
3
4
5
hexaheptaoctanonadeca-
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
=
=
=
=
=
6
7
8
9
10
58
Chapter
3
Molecular Compounds
Molecules
– Electrically neutral particle
– Consists of two or more atoms
Chemical bonds
– Attractions that hold atoms together in molecules
– Arise from sharing electrons between two atoms
– Group of atoms that make up molecule behave
as single particle
Molecular formulas
– Describe composition of molecule
– Specify number of each type of atom present
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
3
Nonmetal Hydrides
Nonmetal hydrides
– Molecule containing nonmetal + hydrogen
– Number of hydrogens that combine with nonmetal =
number of spaces from nonmetal to noble gas in periodic
table
N O F Ne
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
3
Organic Compound Formulas
Molecular formula
– Indicates number of each type of atom in molecule
e.g. C2H6 for ethane or C3H8 for propane
– Order of atoms
• Carbon Hydrogen Other atoms alphabetically
e.g. sucrose is C12H22O11
Emphasize alcohol – write OH group last
– C2H5OH
Structural formula
– Indicate how carbon atoms are connected
– Ethane = CH3CH3
– Propane = CH3CH2CH3
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
3
Nomenclature: Molecular Compounds
• Goal is a name that translates clearly into molecular formula
Naming Binary Molecular Compounds
– Which two elements present?
– How many of each?
Format:
– First element in formula
• Use English name
– Second element
• Use stem and append suffix –ide
– Use Greek number prefixes to specify how many atoms of
each element
Jesperson, Brady, Hyslop. Chemistry:
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Chapter
3
Nomenclature: Binary Molecules
1. hydrogen chloride
1H
2. phosphorous pentachloride
1 Cl
HCl
1P
5Cl
PCl5
3. triselenium dinitride
•
•
•
3 Se 2N
Se3N2
Mono always omitted on first element
Often omitted on second element unless more than one
combination of same two elements
e.g. Carbon monoxide
CO
Carbon dioxide
CO2
When prefix ends in vowel similar to start of element name,
drop prefix vowel
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The Molecular Nature of Matter, 6E
63
Chapter
3
Nomenclature: Exceptions for Binary Molecules
Binary compounds of nonmetals + hydrogen
– No prefixes to be used
– Get number of hydrogens for each nonmetal from
periodic table
– Hydrogen sulfide = H2S
– Hydrogen telluride = H2Te
Molecules with Common Names
– Some molecules have names that predate IUPAC
systematic names
– Water
H2O ▪ Sucrose
C12H22O11
– Ammonia NH3
▪ Phosphine PH3
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
64