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Ch t 1Chapter 1Chemistry: the Central ScienceChemistry: the Central Science
CHEMISTRY , Julia Burge 2nd edition
Copyright McGraw-Hill 2009 1
1 1 The Study of Chemistry1.1 The Study of Chemistry
• Chemistry– the study ofmatter and the changes thatthe study of matter and the changes that matter undergoes
• Matter– anything that has mass and occupies space– anything that has mass and occupies space
1 3 Classification of Matter1.3 Classification of Matter
• Matter is either classified as a substance• Matter is either classified as a substanceor a mixture of substances
• Substance– Can be either an element or a compound– Has a definite (constant) composition and distinct properties
– Examples: sodium chloride, water, oxygen
• States of MatterStates of Matter– Solid
• particles close together in orderly fashion• particles close together in orderly fashion• little freedom of motion• a solid sample does not conform to the shape of p pits container
Li id– Liquid • particles close together but not held rigidly in positionp
• particles are free to move past one another• a liquid sample conforms to the shape of the part f th t i it fillof the container it fills
– Gas• particles randomly spread apart• particles have complete freedom of p pmovement
• a gas sample assumes both shape and volumeg p pof container.
– States of matter can be inter‐converted without changing chemical composition• solid → liquid → gas (add heat) • gas → liquid → solid (remove heat)
States of Matter
Substances• Element: cannot be separated into simpler substances by chemical meanssubstances by chemical means. – Examples: iron, mercury, oxygen, and hydrogen
• Compounds: two or more elements h i ll bi d i d fi i ichemically combined in definite ratios– Cannot be separated by physical means– Examples: salt, water and carbon dioxide
Mixtures • Mixture: physical combination of two or more substances– Substances retain distinct identities– No universal constant compositionNo universal constant composition– Can be separated by physical means
E l / /•Examples: sugar/iron; sugar/water
Molecular Comparison of Substances and Mixtures
f M l l f l tAtoms of an element Molecules of an element
Molecules of a compound Mixture of two elementsMolecules of a compound Mixture of two elements and a compound
• Types of Mixtures– Homogeneous: composition of the mixture is uniform throughout• Example: sugar dissolved in water
– Heterogeneous: composition is not uniform throughout• Example: sugar mixed with iron filings
Classification of Matter
Classify the following
Aluminum foil: substance elementsubstance, element
Baking soda: substance compoundsubstance, compound
Milk:i t hmixture, homogeneous
Air:i hmixture, homogeneous
Copper wire: substance, element
1.3 Scientific Measurement
• Used to measure quantitative properties of mattermatter
• SI (System International) base units
Why are units important?
Mars Climate Orbiter ($125 million) destroyed by wrong conversion units
SI PrefixesSI Prefixes
• Mass: measure of the amount of matter( i h f i i l ll)– (weight refers to gravitational pull)
• Temperature:– Celsius– Celsius
• Represented by °C • Based on freezing point of water as 0°C g pand boiling point of water as 100°C
– KelvinR t d b K ( d i )• Represented by K (no degree sign)
• The absolute scale • Units of Celsius and Kelvin are equal in• Units of Celsius and Kelvin are equal in magnitude
– Fahrenheit (the English system) (°F)
Equations for Temperature ConversionsEquations for Temperature Conversions
59532) F( C oo ×−=
273.15 CK o +=
32C9F oo +×= 32 C5
F +×=
Temperature ConversionsA clock on a local bank reported a t t di f 28oC Wh t i thitemperature reading of 28oC. What is thistemperature on the Kelvin scale?
273 15CK o += 273.15CK +=
o K301273.15C28 K o =+=
Practice
Convert the temperature reading on the local bank (28°C) into the corresponding Fahrenheit temperature.
32 C59F oo +×=5
F8232C289F ooo F82 32C28 59F oo o =+×=
• Volume: meter cubed (m3)Derived unit– Derived unit
– The unit liter (L) is more commonly used in the laboratory setting It is equal to a decimeterlaboratory setting. It is equal to a decimeter cubed (dm3).
• Density: Ratio of mass to volume
– Formula:md =
d d i ( / L)
Vd =
– d = density (g/mL)–m = mass (g)– V = volume (mL or cm3)
(*gas densities are usually expressed in g/L)( gas densities are usually expressed in g/L)
PracticeThe density of a piece of copper wire is 8.96 g/cm3. Calculate the volume in cm3g/of a piece of copper with a mass of 4.28 g.
Vmd =V
30 478g4.28mV 3
3
cm0.478
cmg 8.96
g===
dV
cm
1 4 Properties of Matter1.4 Properties of Matter
• Quantitative: expressed using numbersQuantitative: expressed using numbers• Qualitative: expressed using propertiesPh sical properties b b d d• Physical properties: can be observed and measured without changing the substanceE l l l i i f– Examples: color, melting point, states of matter
• Physical changes: the identity of the substance stays the same– Examples: changes of state (melting, freezing)
• Chemical properties:must be determined• Chemical properties:must be determined by the chemical changes that are observedExamples: flammability acidity corrosiveness– Examples: flammability, acidity, corrosiveness, reactivity
• Chemical changes: after a chemical change, g gthe original substance no longer exists– Examples: combustion, digestion p g
1 5 Uncertainty in Measurement1.5 Uncertainty in Measurement
• Exact: numbers with defined values– Examples: counting numbers, conversionExamples: counting numbers, conversion factors based on definitions
• Inexact: numbers obtained by any method other than countingmethod other than counting– Examples: measured values in the laboratory
Practice
105.5 L + 10.65 L = 116.2 LC l l 116 15 LCalculator answer: 116.15 L Round to: 116.2 L Answer to the tenth position
1.0267 cm x 2.508 cm x 12.599 cm = 32.44 cm3
Calculator answer: 32.4419664 cm3
Round to: 32.44 cm3 round to the smallest number of significant figures
Accuracy and Precision
• Accuracy and precisiony p– Two ways to gauge the quality of a set of measured numbersof measured numbers
–A h l t i– Accuracy: how close a measurement is to the true or accepted value
– Precision: how closely measurements of the same thing are to one another
both accurate and precise
not accurate but precise
neither accurate nor precise
Describe accuracy and precision for each setDescribe accuracy and precision for each set
Student A Student B Student C0 335 g 0 357 g 0 369 g0.335 g 0.357 g 0.369 g0.331 g 0.375 g 0.373 g0.333 g 0.338 g 0.371 gg g g
Average:0.333 g 0.357 g 0.371 g
True mass is 0 370 grams• True mass is 0.370 grams
Student A’s results are precise but not accurate.
Student B’s results are neither precise nor accurate.
Student C’s results are both precise and accurateStudent C s results are both precise and accurate.
1.6 Using Units and Solving Problems
• Conversion factor: a fraction in which the same quantity is expressed one way q y p yin the numerator and another way in the denominator– Example: by definition, 1 inch = 2.54 cm
cm 2.54in 1
in 1cm 2.54
• Dimensional analysis: a problem solving• Dimensional analysis: a problem solving method employing conversion factors to change one measure to another often c a ge o e easu e to a ot e o tecalled the “factor‐label method”
– Example: Convert 12.00 inches to meters• Conversion factors needed: 2.54 cm = 1 in and 100 cm = 1 meter
12 54 m 0.3048cm 100
m1in 1
cm2.54in 12.00 =××
Notes on Problem SolvingNotes on Problem Solving
• Read carefully; find information given and what is asked for
• Find appropriate equations, constants, conversion factors
• Check for sign, units and significant figures• Check for reasonable answer
Practice
The Food and Drug Administration (FDA)d th t di t di i t krecommends that dietary sodium intake
be no more than 2400 mg per day. Whatis this mass in pounds (lb), if 1 lb = 453.6 g?
lb10 5.3g453 6
lb 1mg1000
g 1 mg 2400 3−×=××g 453.6mg1000
Key Points
• Scientific method• Classifying matter• Density• Temperature conversions• Physical vs chemical properties and changes• Physical vs chemical properties and changes • Precision vs accuracy
l l• Dimensional analysis
QUESTIONS?QUESTIONS?
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Chapter 2Atoms, Molecules, and Ions
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2.3 Atomic Number, Mass Number and Isotopes
• The chemical identity of an atom can be determined solely from its atomic number
• Atomic number (Z) ‐ number of protons in the nucleus of each atom of an element– Also indicates number of electrons in the atom—since atoms are neutral
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• Mass number (A) ‐ total number of neutrons and protons present in theneutrons and protons present in the nucleus
• Standard notation:
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• Isotopes• Isotopes – All atoms are not identical (as had been proposed by Dalton)
– Same atomic number (Z) but different ( )mass numbers (A)
• Isotopes of HydrogenIsotopes of Hydrogen– Hydrogen (protium) D t i– Deuterium
– Tritium
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2 4 The Periodic Table2.4 The Periodic Table
• Periods ‐ horizontal rows • Families (Groups) ‐ vertical columns• Families (Groups) vertical columns
– Elements in the same family have similar chemical and physical propertieschemical and physical properties
• Arranged in order of increasing atomic numbernumber
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The Modern Periodic Table
Ti
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• Metals ‐ good conductors of heat and gelectricity (majority of elements on the table, located to the left of the stair step)p)
• Nonmetals ‐ nonconductors (located in• Nonmetals ‐ nonconductors (located in upper right‐hand corner)
• Metalloids ‐ in between metals and nonmetals (those that lie along the separation line)
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Groups (Families) on the Periodic Table
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2.5 The Atomic Mass Scale and A t i MAverage atomic Mass
• Atomic mass is the mass of the atom inAtomic mass is the mass of the atom in atomic mass units (amu)
• Atomic mass unit is defined as a mass exactly equal to one‐twelfth the mass of oneexactly equal to one twelfth the mass of one carbon‐12 atom
• Carbon‐12 (12 amu) provides the standard for measuring the atomic mass of the other
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for measuring the atomic mass of the other elements
2.6 Molecules and Molecular Compounds
• Molecule combination of at least two atoms in a specific arrangement held p gtogether by chemical bonds – May be an element or a compoundMay be an element or a compound– H2, hydrogen gas, is an element H O i d– H2O, water, is a compound
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• Diatomic molecules: – Homonuclear (2 of the same atoms)
• Examples: H2, N2, O2, F2, Cl2, Br2, and I2p 2, 2, 2, 2, 2, 2, 2
– Heteronuclear (2 different atoms)• Examples: CO and HCl
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• Polyatomic molecules: y– Contain more than 2 atoms – Most moleculesMost molecules – May contain more than one element – Examples: ozone O ; white phosphorus P ;– Examples: ozone, O3; white phosphorus, P4; water, H2O, and methane (CH4)
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• Molecular formula – shows exact number of atoms of each element in a molecule– Subscripts indicate number of atoms of each element present in the formula.each element present in the formula.
– Example: C12H22O11
Structural formula ‐ shows the generalStructural formula shows the general arrangement of atoms within the molecule.
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• Naming molecular compoundsg p– Binary Molecular compounds •Composed of two nonmetals•Composed of two nonmetals •Name the first element•Name the second element changing ending to “‐ide” •Use prefixes to indicate number of atoms of each element
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Name the following: NO2
nitrogen dioxideN2O4
dinitrogen tetraoxideg
Write formulas for the following:Write formulas for the following: Diphosphorus pentoxide
P OP2O5Sulfur hexafluoride
SF53
SF6
Common NamesCommon Names
B2H6 diboraneSiH4 silaneSiH4 silaneNH3 ammoniaPH h hiPH3 phosphineH2O waterH2S hydrogen sulfide
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• Acid ‐ a substance that produces phydrogen ions (H+) when dissolved in water
• Binary acids: Many have 2 names– Many have 2 names •Pure substance•Aqueous solution
– Example: HCl, hydrogen chloride, when p , y g ,dissolved in water it is called hydrochloric acid
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y
• Naming binary acids– Remove the “–gen” ending from g ghydrogen (leaving hydro–)
– Change the “–ide” ending on the secondChange the ide ending on the second element to “–ic”Combine the two words and add the– Combine the two words and add the word “acid.”
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Name the following: HBrhydrogen bromide hydrobromic acidhydrogen bromide hydrobromic acidH2Sh d lfid h d lf i idhydrogen sulfide hydrosulfuric acidWrite formulas for the following: Hydrochloric acid
HCl(aq)HCl(aq)Hydrofluoric acid
HF( )57
HF(aq)
• Organic compounds ‐ contain carbon g pand hydrogen (sometimes with oxygen, nitrogen, sulfur and the halogens.)g g )– Hydrocarbons ‐ contain only carbon and hydrogen
– Alkanes ‐ simplest examples of hydrocarbons
– Many derivatives of alkanes are derived by replacing a hydrogen with one of the f i lfunctional groups. • Functional group determines chemical
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properties
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• Empirical formulas reveal the elements p fpresent and in what whole‐number ratio they are combined.y
Molecular(explicit) Empirical(simplest)Molecular(explicit) Empirical(simplest)H2O2 HON H NHN2H4 NH2
H2O H2O
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2.7 Ions and Ionic Compounds
• Ion ‐ an atom or group of atoms that has a net positive or negative charge
• Monatomic ion ‐ one atom with a positive or negative charge
• Cation ion with a net positive charge due to the loss of one or more electrons
• Anion ion with a net negative charge due the gain of one or more electrons
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Common Monatomic Ions
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• Naming ions– Cations from A group metals
• Name the element and add the word “ion”E l N d• Example: Na+, sodium ion
– Cations from transition metals with some exceptionsexceptions• Name element • Indicate charge of metal with Roman• Indicate charge of metal with Roman numeral
• Add word “ion”• Example: Cu2+ ,copper(II) ion
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– Anions N h l d dif h di• Name the element and modify the ending to “‐ide” E l Cl hl id• Example: Cl−, chloride
Polyatomic ions ions that are a combination• Polyatomic ions ‐ ions that are a combination of two or more atomsNotice similarities number of oxygen atoms– Notice similarities ‐ number of oxygen atoms and endings for oxoanions• Nitrate NO − and nitrite NO −• Nitrate, NO3 and nitrite, NO2
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• Ionic compounds ‐ represented by empirical formulas– Compound formed is electrically p yneutral
– Sum of the charges on the cation(s) andSum of the charges on the cation(s) and anion(s) in each formula unit must be zerozero
– Examples: Al3+ and O2− Al OAl and O Al2O3Ca2+ and PO43− Ca3(PO4)2
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Formation of an Ionic Compound
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Write empirical formulas for p• Aluminum and bromide
AlBrAlBr3• barium and phosphate
Ba3(PO4)2• Magnesium and nitrate
Mg(NO3)2• Ammonium and sulfate• Ammonium and sulfate
(NH4)2SO4
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• Naming ionic compounds• Naming ionic compounds– Name the cation Na e the a io– Name the anion
– Check the name of cation • If it is a A group metal you are finished f l h• If it is a transition metal, with some exceptions, add the appropriate Roman numeral to indicate theRoman numeral to indicate the positive ionic charge
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Write names for the following: • KMnO4
potassium permanganatep p g• Sr3(PO4)2
strontium phosphatestrontium phosphate• Co(NO3)2
cobalt(II) nitrate • FeSO44
iron(II) sulfate
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• Hydrates ‐ compounds that have a specific number of water molecules within their solid structure– Hydrated compounds may be heated to remove the water forming an anhydrous
dcompound– Name the compound and add the word h d t I di t th b f thydrate. Indicate the number of water molecules with a prefix on hydrate. Example: CuSO 5 H O• Example: CuSO4 · 5 H2O –Copper (II) sulfate pentahydrate
74See picture…
CuSO4
CuSO4.5H2O
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Key Points• Isotopes
• Periodic table; families and periods; metals• Periodic table; families and periods; metals, nonmetals and metalloids
• Average atomic mass
Naming and writing formulas for• Naming and writing formulas for– Binary molecular compounds
d– Binary acids– Ionic compounds
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– Hydrates
QUESTIONS?QUESTIONS?
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