chapter 11: chemical elements

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Chapter 11: Chemical Elements

Homework: All questions on the “Multiple-Choice” and the odd-numbered questions on“Exercises” sections at the end of the chapter.

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Chemistry

• Chemistry – a division of PhysicalScience

• Chemistry – deals with the compositionand structure of matter and thereactions by which substances arechanged into other substances

• Egyptian, Chinese, and Mesopotamians– wine making, worked metals, dyes,glass, pottery, embalming fluids (asearly as 3500 BC

Intro


Alchemy & Modern Chemistry

• Flourished from 500 – 1600 AD

• Main objectives (never reached):

– Change common metals to gold

– Find an “elixir of life”

• Modern Chemistry -- began in 1774,with the Frenchman Antoine Lavosier –quantitative methods and avoidedmysticism, superstition, and secrecy

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Major Divisions of Chemistry

• Physical Chemistry – applies thetheories of physics

• Analytical Chemistry – identifies whatand how much is present

• Organic Chemistry – carboncompounds

• Inorganic Chemistry – non-carboncompounds

• Biochemistry – chemical reactions thatoccur in living organisms

Intro


88 Naturally Occurring Elements

• Either singly or in chemical combination– the 88 naturally occurring elementscomprise virtually all matter.

• Their chemical and physical propertiesaffect us continually

• This chapter (#11) –– Classification of matter

– Discuss elements

– Discuss the Periodic Chart

– Discuss the Naming of Compounds

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Classification of Matter

• In Chapter 5 we saw that matter can beclassified by its physical phase or state –solid, liquid, and gas

• Matter – anything that has mass

• Chemists use this classification, but alsodivide matter into several other classifications

• Pure Substance – element or compound

• Mixture – homogeneous or heterogeneous

Section 11.1

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Pure Substance

• Pure Substance – type of matter in which allsamples have fixed composition and identicalproperties

– Element – all atoms have same # of protons (gold,sulfur, oxygen)

– Compound – two or more elements chemicallycombined in a definite, fixed ratio by mass (puresalt, topaz crystal, distilled water)

• A compound can be broken into its separatecomponents only by chemical processes(electrical current)

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Natural minerals -good examples of Compoundsof Pure Substances

Faceted TopazAl2SiO4(OH,F)2

Halite NaCl “rock salt”

Rhodochrosite MnCO3

Copyright © Bobby H. Bammel. All rights reserved.

Section 11.1

“Lone Star Cut”


Compound vs. Component ElementsCompound - usually different from the

components


Mixture

• Mixture – type of matter composed of varyingproportions of two or more substances thatare only physically mixed and not chemicallycombined

– Homogeneous (a solution)– uniform throughout(coffee, alloy) technically, it should bemixed/uniform on the atomic level

– Heterogeneous – non-uniform (pizza, oil/water), atleast two components can be observed

• Formed and broken down by physicalprocesses (dissolving, evaporation)

Section 11.1


Chemical Classification of Matter


Liquid Solutions

• Solvent – the liquid or the substance in the larger quantity

• Solute – the substance dissolved in the solvent

Section 11.1

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Aqueous Solutions

• Aqueous Solution (aq)– a solution inwhich water is the solvent– When dissolved & stirred the distribution of

the solute is the same throughout(homogeneous)

• Unsaturated Solution – more solute canbe dissolved in the solution at the sametemp.

• Saturated Solution – maximum amountof solute is dissolved in the solvent


Saturated Solution

• A dynamicequilibriumexistsbetween thesolutedissolving andthe solutecrystallizing

Section 11.1


Solubility

• Solubility – the amount of solute that willdissolve in a specified volume or massof solvent (at a given temperature) toproduce a saturated solution

• If the temperature is raised thesolubilities for most solids increase


The Effect of Temperature on Solubilities ofSalts in Water

• Usually hotter waterwill dissolve moresolute

Section 11.1


Supersaturated Solutions

• When unsaturated solutions areprepared at high temperatures and thencooled, the saturation point may bereached as the solution cools

• However, is no crystals are present,crystallization may not take place

• Result Supersaturated Solution –contains more than the normalmaximum amount of dissolved solute atthe given temperature


Seed Crystal added to a SupersaturatedSolution

Section 11.1

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Solubility of Gases

• The solubility of gases increases withincreasing pressure– Example: manufacture of soft drinks, CO2 is forced

into the beverage at high pressure

• Once the soft drink is opened, the pressureinside the container is reduced to normalatmospheric pressure and the CO2 startsescaping

• The solubility of gases decreases withincreasing temperature (hot soft drinks quicklylose their CO2)


Robert Boyle

• In 1661 Robert Boyle proposed that thedesignation element be applied only tosubstances that could not be separated intocomponents by any method

• In addition Boyle initiated the practice ofcarefully and completely describingexperiments so that anyone might repeat andconfirm them

– Due to this procedure (carefully documentingexperiments) scientists have been able to build onprevious knowledge

Section 11.2


Discovery of the Elements

• The earliest civilizations isolated 12 elements;gold, silver, lead, copper, tin, iron, carbon,sulfur, antimony, arsenic, bismuth, andmercury – later all 12 proved to be elements

• Phosphorus was isolated (from urine) in 1669

– P is the first element whose date of discovery isknown

• By 1746, platinum, cobalt, and zinc had allbeen discovered


Discovery of the Elements

• Around 1808 Davy, an English Chemist, usedelectricity from the recently invented battery tobreak down compounds, thereby isolating sixadditional elements (Na, K, Mg, Ca, Ba, Sr)

• By 1895 a total of 73 element were known

• During the next three years the noble gasesHe, Ne, Kr, and Xe were discovered

• In addition to the naturally occurring elements,26 synthetic elements have now been created

Section 11.2


Discovery of the ElementsShowing how many elements were known at points in

history


Occurrence of the Elements

• Human Body = 65% oxygen & 18% carbon

• Analyses of electromagnetic radiation fromspace indicates that the universe consists of:

– Hydrogen – 75% (simplest element)

– Helium – 24% (second most simple element)

– Others – 1%

• Earth’s Atmosphere = 78% nitrogen, 21%oxygen, and about 1% argon

• Earth’s Core = 85% iron & 15% nickel

Section 11.3

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Relative Abundance (by Mass) of Elementsin Earth’s Crust

Note that 74% ofthe mass of theEarth’s crust iscomposed of onlytwo elements –oxygen & silicon


Atoms

a) The individual units (atoms) packed in arepeating pattern

b) Noble gases that occur as single atoms

c) Diatomic atoms (hydrogen)

Section 11.3


Molecules

• Molecule – an electrically neutral particlecomposed of two or more atomschemically combined

• If the atoms are that same element, thenthe molecule is of an element

– Element examples: H2 or N2

• If the atoms are different elements, thanethe molecule is of a compound

– Compound examples: H2O or NH3


Representations of Molecules

Section 11.3


Seven common elements that exist asdiatomic molecules

• These atoms are tooreactive to exist asindependent atoms.

• When writing formulas w/these seven elements weus the diatomic form: H2

+Cl2 2HCl


Allotropes

• Allotrope – two or more forms of thesame element that have differentbonding structures in the same physicalphase

• Example: Diamond and Graphite

• Both pure Diamond and pure Graphiteare each 100% carbon (C), and areboth solid

• But the atomic arrangement of thecarbon atoms is different

Section 11.3

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Three Allotropes of CarbonDiamond, Graphite, and Buchminsterfullerene (C60)


Oxygen also has two AllotropesOxygen gas (O2) and Ozone (O3)

Section 11.3


The Periodic Table

• The periodic table puts the elements inorder of increasing atomic number, intoseven horizontal rows, called periods

• The elements’ properties show regulartrends going up or down these periods

• In 1869 the Russian Chemist,Mendeleev, published the originalperiodic table

• The fifteen vertical columns in theperiodic table are called groups


The Periodic TableDivided into 7 Periods and 18 Groups

Section 11.4


One way to Classify the periodic table

RepresentativeElements (green)

TransitionalElements (blue)

Inner TransitionElements (purple)


Metals & Nonmetals

• Another way of classifying the elements isinto – metals and nonmetals

• A metal is an element whose atoms tend tolose electrons during chemical reactions (+)

• A nonmetal is an element whose atoms tendto gain (or share) electrons (-)

• The metallic character of the elementsincreases as one goes down a group, anddecreases across (left to right) a period

Section 11.4

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Metals/Nonmetals in the Periodic Table


General Properties of Metals & Nonmetals

Section 11.4


Electron Configuration

• Electrons are located in energy levels orshells that surround the nucleus

– Level 1 – maximum of 2 electrons



• The chemical reactivity of the elementsdepends on the order of electrons inthese energy levels


We will generally only discuss the“representative elements” – Groups 1A - 8A

Section 11.4


Valence Electrons

• The outer shell of an atom is known asthe valence shell

• The electrons in the outer shell arecalled the valence electrons

• The valence electrons are the electronsinvolved in forming chemical bonds – sothey are extremely important

• Elements in a given group all have thesame number of valence electrons (and\ similar chemical properties)


Guidelines: Shell Electron Configurations

• The number of electrons in an atom is thesame as the element’s atomic number (Z)

• The number of shells that contain electronswill be the same as the period number that itis in

• For the A group (representative) elements,the number of valence electrons is the sameas the group number

• Let’s look at Example 11.1 & ConfidenceExercise 11.1 on pages 294-295 in your text

Section 11.4

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Electron Shell Distribution for Periods 1,2,3This is an important and very useful figure!

Protons (Z)Electrons

Section 11.4


The Periodic Nature of Atomic Size

• The atomic size of the elements also variesperiodically (refer to the Periodic Table) – from0.074 nm (H) to 0.47 nm (Cs)

• Atomic size increases down a group

• Atomic size decreases across a period

• The atoms on the far left are the largest due toless charge (fewer protons) in the nucleus and\ the outer electrons are more loosely bound


Relative Atomic Sizes

Note - the Periodic Table can be used to determine relative atomic size

Section 11.5


Ionization Energy – also Periodic

• Ionization energy – the amount ofenergy that it takes to remove anelectron from an atom

• Ionization energy increases across aperiod due to additional protons in thenucleus

• Ionization energy decreases down agroup because of the additional shellssituated between the nucleus and theouter electron shell.


Ionization Energy Trend

Section 11.5

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Chemical Formulas

• In order to easily and conveniently discusschemistry we can use their chemical formulas

• Chemical formulas are written by putting theelements’ symbols adjacent to each other –usually w/ the more metallic element first

• A subscript following each symbol designatesthe number of atoms – H2O

• Some compounds have special names


ElevenCompoundswith Special

Names

Section 11.5


Naming a Binary Compound(Metal + Nonmetal)

• Binary => two-element compound

• First give the name of the metal and thegive the name of the nonmetal,changing its ending to – “ide”

• NaCl sodium chloride

• Al2O3 aluminum oxide

• Ca3N2 calcium nitride


“ide” Nomenclature

Section 11.5


Compounds of Two Nonmetals

• The more metallic or less nonmetallicelement (farther left or farther downperiodic chart) is usually written first inthe formula and named first

• The second element is named using the“ide” ending

• Greek prefixes are used to designatethe number of atoms in the molecule


Examples:

• HCl hydrogen chloride

• CS2 carbon disulfide

• PBr3 phosphorus tribromide

• IF7 iodine heptafluoride

Section 11.5

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Polyatomic Ions

• Ion – an atom or chemical combination ofatoms having a net electric charge

• Monatomic ion – an ion formed from a singleatom (Cl-)

• Polyatomic ion – an electrically chargedcombination of atoms (CO3

2-)

• Name the metal and then the polyatomic ion

– ZnSO4 zinc sulfate

– NaC2H3O2 sodium acetate

– Mg(NO3)2 magnesium nitrate

– K3PO4 potassium phosphate


CommonPolyatomic Ions

Section 11.5


Acids -- when hydrogen (H) is combinedwith polyatomic ions


Naming Compounds - Example

• H2SO4 sulfuric acid (special name)

• ZnCO3 zinc carbonate (metal + polyatomicion)

• Na2S sodium sulfide (binary compound ofmetal + nonmetal)

• NH3 ammonia (special name)

• NH4NO3 ammonium nitrate (ammoniumion + polyatomic ion)

Section 11.5


Flowchart: Naming Compounds


Groups of Elements

• Recall that in the Periodic Table eachindividual column is called a group

• All the elements in a group have thesame number of valence electrons

• If one element in a group reacts with asubstance – the other elements in thegroup usually react similarly

• The formulas of the compounds createdare also similar

• We will discuss four of these groups …

Section 11.6

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Alkali Metals Noble Gases

Alkaline Earth Metals Halogens




1 2 7 8

Section 11.6


Noble (Inert) Gases – Group 8A

• The exist as single atoms (monatomic)

• Almost never react and form compounds

• Noble gases have 8 electrons in their outershells (except He that has a full shell with 2)– Eight electrons in the outer shell must be VERY

stable

• “Neon” signs contain minute amounts ofvarious noble gases – electric current glow!

• Argon gas is used inside light bulbs becauseeven at high temps. it will not react with thetungsten filament (W)




1 2 7 8

Section 11.6


Alkali Metals – Group 1A (not H)

• Each alkali metal atom has only one valenceelectron

• \ tends to lose this electron ( +) and readilyreact with other elements – active metals

• Na & K are abundant (Li, Rb, Cs are rare)

• So reactive w/ oxygen and water that they mustbe stored in oil

• NaCl, K2CO3 (potash), Na2CO3 (washing soda).NaOH (lye), NaHCO3 (baking soda)

• Predict formulas KCl, LiCO3




1 2 7 8

Section 11.6

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Halogens – Group 7A

• Each halogen atom has seven valence electrons

• \ tends to gain an electron and readily reactwith other elements active nonmetals

• Only occur in nature as a compound, but whenpurified occur as a diatomic molecule (F2, Cl2) –generally poisonous

• F is the most reactive – will corrode Pt andcause wood, rubber, water catch fire on contact

• Iodine is necessary for proper thyroid function

• AlCl3 (aluminum chloride), NH4F (ammoniumfluoride), CaBr2 (calcium bromide)


Iodine Deficiency Goiter

• A lack ofiodine in thediet can leadto anenlargedthyroid gland

Section 11.6




1 2 7 8


Alkaline Earth Metals – Group 2A

• This group contains two valence electrons, andtend to lose two electrons ( +2)

• Not as chemically active as alkali metals (1A),and are generally harder and stronger

• Be2Al2(SiO3)6 – (beryl) , Mg(OH)2 (milk ofmagnesia), CaCO3 (calcite), Ca3(PO4)2 (bones& teeth), BaSO4 (barite); Sr (red) & Ba (green)give color in fireworks

• Ra is radioactive – RaCl2 used on watch dials(glowed in dark) until a number of Swiss dial-painters came down with stomach cancer!!

Section 11.6


Hydrogen – Group 1A usuallysometimes Group 7A

• Although a nonmetal, H usually reacts likea alkali metal (HCl, H2S)

• Sometimes reacts like a halogen – NaH,CaH2

• At room temp. – colorless, odorless,diatomic

• Lightest element – was used in earlydirigibles

• Will burn in air to form water


Hindenburg, Lakehurst, NJ; 5/6/37

• UsedflammableH forbuoyancy.Airshipstoday useHe.

Section 11.6

chapter 11: chemical elements

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