Overview of the Basics
CHAPTER 1-3 Review
Chemistry: The Molecular Nature of Matter, 6th editionBy Jesperson, Brady, & Hyslop
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CHAPTER 1-3 Review
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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• Empirical fact• Something we see, hear, taste, feel, or smell• Something we can measure • Organize data so we can see relationships
2. Law or Scientific Law • Usually an equation• Based on results of many experiments• Only states what happens• Does not explain why they happen
3. Hypothesis• Mental picture that explains observed laws • Tentative explanation of data• Make predictions• Leads to further tests• Go to laboratory and perform experiments
4. 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
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Chapter 1 Elements
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
• 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 sulfur
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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Chapter 1 Elements
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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Chapter 1 Classification of Matter
http://ridenourmhs.wikispaces.com/ESUnit2
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Chapter 2 Chemical vs Physical Properties
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Physical properties Can be observed without changing chemical makeup of substance
Chemical properties
Solids: Fixed shape and volumeParticles are close together
Liquids: Fixed volume, but take container shapeParticles are close together
Gases: Expand to fill entire containerParticles separated by lots of space
• Chemical change or reaction that substance undergoes
• Chemicals interact to form entirely differentsubstances with different chemical and physical properties
• Describe behavior of matter that leads to formation of new substance
• “Reactivity" of substancee.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. WaterMass 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 trioxideMass S 32.06 g 32.06 g
Mass O 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: The Molecular Nature of Matter, 6E
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Chapter 1 Hydrates
• Crystals that contain water moleculese.g. Plaster: CaSO4∙2H2O calcium sulfate dihydrate
– Water is not tightly held• Dehydration
– Removal of water by heating– Remaining solid is anhydrous (without water)
Blue =
CuSO4 •5H2O
White = CuSO4
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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Chapter 1 Depicting Molecules
H C H
H
H
CH4
methane
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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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: The Molecular Nature of Matter, 6E
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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 CO2 + 2H2O 4H + 4O + C = 4H + 4O + C
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Chapter 1 Balanced Chemical Equations
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Ex. 2C4H10 + 13O2 8CO2 + 10H2O
4 C and 10 H per molecule 2 O per
molecule
2 H and 1 O per
molecule
1 C and 2 O per
moleculeEx. 2C4H10 + 13O2 8CO2 + 10H2O
2 molecules of C4H10
13 molecules of O2
10 molecules of C4H108
molecules of CO2
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Chapter 2 Intensive vs Extensive Properties
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Intensive properties – Independent of sample
size– Used to identify
substancese.g. Color
Density Boiling point Melting point Chemical reactivity
Extensive properties – Depend on sample sizee.g. volume and mass
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Chapter 2 Measurements
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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
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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: The Molecular Nature of Matter, 6E
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Chapter 2
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
10001
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 =
g
4. Temperature (T)273.15 K = 0°C
4 Common Lab Measurements
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Chapter 2
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Uncertainty in Measurements
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
Measurements all inexactLimitations of reading instrument
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Chapter 2 Significant Figures
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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
Scientific convention: All digits in measurement up to and including first estimated digit are significant.
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Chapter 2 Significant Figures
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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 significante.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.
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Chapter 2 Significant Figures: Rounding
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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 6.6b. Odd, it rounds up.e.g. 6.35 is rounded to 6.4
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Chapter 2 Significant Figures: Calculations
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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• Answer has same number
of decimal places as quantity with fewest number of decimal places.
12.9753 319.5+ 4.398
4 decimal places1 decimal place3 decimal places1 decimal place336.9
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Chapter 2 Significant Figures: Exact Numbers
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
• 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
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Chapter 2 Scientific Notation
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
• 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
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Chapter 2 Accuracy & Precision
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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
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Chapter 2 Dimensional Analysis
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
• Also called the Factor Label Method• Not all calculations use specific equation• Use units (dimensions) to analyze problemConversion Factor • Fraction formed from valid equality or
equivalence between units• Used to switch from one system of
measurement and units to another
Given Quantity
Desired Quantity
Conversion Factor
× =
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Chapter 2 Dimensional Analysis
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Example: Convert 0.097 m to mm.• Relationship is 1 mm = 1 × 10–3 m• Can make two conversion factors
• Since going from m to mm use one on left.m 101
mm 13 mm 1
m 101 3
m 101mm 1m 097.0 3
= 97 cm
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Chapter 2 Density
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
• Ratio of object’s mass to its volume
• 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
volumemassdensity
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Chapter 3
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Rutherford Nuclear Atom
Discovery of the NucleusRutherford
Alpha scattering expt
Discovery of Protons1918 Rutherford
Mass spectrometer
Discovery of electron mass and charge
Millikan Oil Drop expt
Discovery of the Electron1897 ThomsonCathode ray tube expt
Discovery of Neutron:1932 Chadwick
Discovery of Subatomic Particles in the late 1800’s and early 1900s
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Chapter 3 Properties of Subatomic Particles
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Particle Mass (g) Electrical Charge Symbol
Electron
9.10939 10–28 –1
Proton 1.67264 10–24 +1
Neutron 1.67495 10–24 0
e01
p11
11 ,H
n10
Three kinds of subatomic particles of principal interest to chemists
Nucleus (protons + neutrons)
Electrons
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Chapter 3 Atomic Notation
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Atomic number (Z) = Number of protons that atom has in nucleusIsotopes = Atoms of same element with different
masses– Same number of protons ( )– Different number of neutrons ( )
Isotope Mass number (A)– A = (number of protons)+(number of neutrons) =
Z + N– For charge neutrality, number of electrons and
protons must be equalAtomic Symbols = Summarize information about
subatomic particles– Every isotope defined by two numbers Z and A
Ex. What is the atomic symbol for helium? He has 2 e–, 2 n and 2 p Z = 2, A = 4
p11
n01
He42
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Chapter 3 Isotopes
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
• 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)
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Chapter 3 Carbon-12 Atomic Mass Scale
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
• Need uniform mass scale for atomsAtomic 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
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Chapter 3 Calculating Atomic Mass
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
• Generally, elements are mixtures of isotopese.g. HydrogenIsotope Mass % Abundance 1H 1.007825 u 99.985 2H 2.0140 u 0.015How do we define atomic mass?
– Average of masses of all stable isotopes of given elementHow do we calculate average atomic mass?
– Weighted average– Use isotopic abundances and isotopic masses
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Chapter 3 Periodic Table
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
– Summarizes periodic properties of elementsEarly 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
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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
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
1A 2A B B 7A 8A
Alkali Metals Alkaline Earth
Metals
Transition Metals
Lanthanide & Actinide
Halogens Nobel Gases
Very reactive
Metals except for H
+1 ions
React with Oxygen to form compounds that dissolve into alkaline solutions in water
Reactive
+2 ions
Oxygen compounds are strongly alkaline
Many are not water soluble
Metals
Form ions with several different charges (oxidation states)
Tend to form +2 and +3 ions
Lanthanides 58 – 71
Actinides 90 – 103
Actinides are radioactive
Reactive
Form diatomic molecules in elemental state
-1 ions
Salts with alkali metals
Inert
Heavier elements have limited reactivity
Do not form ions
Monoatomic gases
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Chapter 3 Metals, Nonmetals, and Metalloids
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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Chapter 3 Metals, Nonmetals, and Metalloids
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Metals Nonmetals Metalloids
• Metallic luster, malleable, ductile, hardness variable
• Conduct heat and electricity
• Solids at room temperature with the exception of Hg
• Chemical reactivity varies greatly: Au, Pt unreactive while Na, K very reactive
• Brittle
• Insulators, non-conductors of electricity and heat
• Chemical reactivity varies
• Exist mostly as compounds rather then pure elements
• Many are gases, some are solids at room temp, only Br2 is a liquid.
• Metallic shine but brittle
• Semiconductors: conduct electricity but not as well as metals: examples are silicon and germanium
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Chapter 3 Ions and Ionic Compounds
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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
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Chapter 3 Ions and Ionic Compounds
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Metal + Non-metal ionic compound 2Na(s) + Cl2(g) 2NaCl(s)
Na+ + ClNa + Cl NaCl(s)
e
Michael Watson Richard Megna/Fundamental Photographs Richard Megna/Fundamental Photographs
Anions = Negatively charged ions Cations = Positively charged ions
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Chapter 3 Ions and Ionic Compounds
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Electrical conductivity requires charge movementIonic 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
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Chapter 3 Ions and Ionic Compounds
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Negative (–) charge on anion = number of spaces you have to move to
right to get to noble gas
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Chapter 3 Rules for Writing Ionic Formulas
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
1. Cation given first in formula2. Subscripts in formula must produce electrically neutral
formula unit3. 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
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Chapter 3 Determining Ionic Formulas
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
“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?
Al2O3Al3+ O2–
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Chapter 3 Transition Metal and Post-Transition Metal Ions
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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Chapter 3 PolyatomicIons
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Example: What is the formula of the ionic compound formed between ammonium and phosphate ions?
Ammonium = NH4+
Phosphate = PO43–
(NH4)+ (PO4)3–
(NH4)3PO4
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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
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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+ chromium(III)
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Chapter 3 Nomenclature: Ionic Compounds
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Anions:– Monatomic anions named by adding
“–ide” suffix to stem name for element
– Polyatomic ions use names in Table 3.5
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Chapter 3 Nomenclature: Hydrates
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
• 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
mono- = 1 hexa- = 6di- = 2 hepta- = 7tri- = 3 octa- = 8tetra- = 4 nona- = 9penta- = 5 deca- = 10
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Chapter 3 Molecular Compounds
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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 particleMolecular formulas
– Describe composition of molecule– Specify number of each type of atom present
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Chapter 3 Nonmetal Hydrides
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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
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Chapter 3 Organic Compound Formulas
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Molecular formula – Indicates number of each type of atom in moleculee.g. C2H6 for ethane or C3H8 for propane– Order of atoms
• Carbon Hydrogen Other atoms alphabeticallye.g. sucrose is C12H22O11
Emphasize alcohol – write OH group last– C2H5OH
Structural formula– Indicate how carbon atoms are connected– Ethane = CH3CH3
– Propane = CH3CH2CH3
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Chapter 3 Nomenclature: Molecular Compounds
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
• Goal is a name that translates clearly into molecular formulaNaming 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
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Chapter 3 Nomenclature: Binary Molecules
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
1. hydrogen chloride
2. phosphorous pentachloride
3. triselenium dinitride
• Mono always omitted on first element• Often omitted on second element unless more than one
combination of same two elementse.g. Carbon monoxide CO Carbon dioxide CO2
• When prefix ends in vowel similar to start of element name, drop prefix vowel
1 H 1 Cl HCl
1 P 5Cl PCl5
3 Se 2N Se3N2
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Chapter 3 Nomenclature: Exceptions for Binary Molecules
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
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
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