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Chapter 5: Atomic Structureand the Periodic Table

• Atoms• Structure of the Nuclear

Atom• Distinguishing Between

Atoms• The Periodic Table:

Organizing the Elements

Chapter 5:Atomic Structure and the

Periodic Table

Atoms

Chapter 5: Atomic Structure and the Periodic Table

-- Atoms --Early Models of the Atom• Democritus

– Suggested the existence of atoms– Thought that atoms were indivisible and

indestructible– Problems

• Didn’t explain chemical behavior• No experimental support


-- Atoms --

• Dalton’s Model– Studied ratios in which

elements combine in chemical reactions

– Dalton’s Atomic Theory• All elements are made of tiny indivisible particles

called atoms.• Atoms of the same element are the same but are

different from those of other elements.• Atoms of different elements can physically mix or

can chemically combine with each other in simple, whole-number ratios to form compounds.


-- Atoms --• Chemical reactions occur when atoms are

separated, joined, or rearranged. Atoms of one element never change into other kinds of atoms in chemical reactions.

– Inaccuracies in Dalton’s Theory• Atoms are divisible.• Atoms of the same element aren’t entirely the

same always.• Current theory of the atom

– Smallest particle of an element that still has all its properties

– Extremely small and typically measured in angstroms (1 x 10-10 m)


Periodic Table

Structure of the Nuclear Atom


-- Structure of the Nuclear Atom --• Electrons

– Negatively charged subatomic particle– Extremely light compared to other subatomic

particles– Experiments

• J.J. Thompson–Use of the cathode ray tube–Attraction and repulsion–Determined the mass of an electron is about

1/2000 the mass of a hydrogen atom• Robert A. Millikan


-- Structure of the Nuclear Atom --–Calculated an accurate value for the mass of an

electron–Discovered the electron carries exactly one unit

of negative charge–Mass of an electron is 1/1840 of a hydrogen

atom• Protons and Neutrons

– Properties of atoms and charge• Atoms have no net electric charge; they are neutral• Electric charges are carried by matter• Electric charges always exist in whole-number

multiples of a single basic unit


-- Structure of the Nuclear Atom --• Charges have to cancel each other out to be

neutral– Protons

• Positively charged subatomic particles• Discovered by E. Goldstein in cathode ray tube• 1840 times more massive than an electron• Proton carries exactly one unit of positive charge

– Neutrons• Neutrally charged subatomic particles• Discovered by James Chadwick• As massive as a proton


-- Structure of the Nuclear Atom --• The Atomic Nucleus

– Original theory of uniform distribution– Ernest Rutherford

• Alpha particles (helium atoms that lost two electrons) fired against gold foil

• Atoms have a lot of empty space• Core of atom is the nucleus composed of protons

and neutrons


-- Structure of the Nuclear Atom --


-- Structure of the Nuclear Atom --


-- Structure of the Nuclear Atom -- Chapter 5:Atomic Structure and the

Periodic Table

Distinguishing Between Atoms


-- Distinguishing Between Atoms --Atomic Number• Number of protons in an atom• Same as the number of electrons in a neutral atom

(since protons cancel out electron charges in a neutral atom)

Mass Number• Total number of protons and neutrons in an atom• Mass of an atom does not come from electrons

because of how light they are• # of neutrons = Mass number - atomic number


-- Distinguishing Between Atoms --

Also known as oxygen-16 (“element” - mass number)Isotopes• Atoms with the same number of protons but different numbers of

neutrons• Have different mass numbers because of the same number of

protons and electrons but differing neutrons• Chemically alike because of protons and electrons



Isotopes of Hydrogen



1) Mass number?2) Atomic number?3) p+ (protons)?4) n0 (neutrons)?5) e- (electrons)?

Atomic Number

Mass Number OR Copper-63

63 (always given)29 (given and on periodic table)29 (definition of atomic number)34 (mass number - atomic number)29 (has to cancel out 29 protons)




Uranium-234234 (always given)92 (on periodic table)92 (definition of atomic number)142 (mass number - atomic number)92 (has to cancel out 92 protons)




80 (always given)35 (given and on periodic table)35 (definition of atomic number)45 (mass number - atomic number)36 (extra 1-, so add to gain an

electron since it is negative)




40 (always given)20 (given and on periodic table)20 (definition of atomic number)20 (mass number - atomic number)18 (positive 2, so lost 2 electrons

since it is positive)


-- Distinguishing Between Atoms --Atomic Mass• Possible to predict using a mass spectrometer• Atomic mass units

– Abbreviated “amu”– 1/12 the mass of a carbon-12 atom

• Relative abundance– Each isotope of an element has a fixed mass and

occurs in a certain percentage in nature, like how nitrogen makes up a certain percentage of air, oxygen another percentage, etc.


-- Distinguishing Between Atoms --Natural Percent Abundance of Oxygen



Average atomic mass (“weighted average”):= 0.99759 15.995 + 0.00037 16.995 + 0.00204 17.999= 15.956 + 0.0063 + 0.0367= 15.999 amu


-- Distinguishing Between Atoms --• Atomic mass inferences

– Average atomic masses are never whole numbers– Example: average atomic mass of copper

• Average atomic mass of copper is 63.546 among isotopes of copper-63 and copper-65

• Since 63 is closer to 63.546 than 65, copper-63 must be more abundant than copper-65


Periodic Table

The Periodic Table:Organizing the Elements


-- The Periodic Table: Organizing the Elements --Development of the Periodic Table• Dmitri Mendeleev

– Listed the elements in order of increasing atomic mass

– Constructed the first periodic table, an arrangement of elements according to similarities in their properties

– Created by establishing patterns in physical and chemical properties

• Henry Moseley– Determined the atomic number of the atoms of the

elements– Created a periodic table based on atomic number as

is currently used today


-- The Periodic Table: Organizing the Elements --The Periodic Table


-- The Periodic Table: Organizing the Elements --The Modern Periodic Table• Elements are listed in order of increasing atomic

number• Periods

– Horizontal rows– Patterns of properties within a period repeat from

period to period– Periodic Law

• When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.


-- The Periodic Table: Organizing the Elements --• Groups

– Vertical rows– Also called “families”– Identified with a number or letter– Representative elements

• Wide range of physical and chemical properties• Group A elements

– Transition metals and inner transition metals• More similar physical and chemical properties• Transition metals: Group B elements• Inner transition metals: lanthanide and actinide

series


-- The Periodic Table: Organizing the Elements --– Alkali metals

• Group 1A elements• Highly reactive• Examples: sodium (Na), potassium (K)

– Alkaline earth metals• Group 2A elements• Not as reactive as alkali metals• Examples: barium (Ba), calcium (Ca)

– Halogens• Group 7A elements• Very reactive - do not exist as single atoms• Examples: fluorine (F), chlorine (Cl)


-- The Periodic Table: Organizing the Elements --– Noble gases

• Group 8A elements• Do not react at all except in extreme cases• Examples: Helium (He), Neon (Ne)

• Metals vs. Nonmetals– Metals

• On the left side of the periodic table• Conduct electricity well and have a luster• Make up a majority of the elements on the periodic

table• Exist primarily as solids at room temperature (with

the exception of mercury which is a liquid at room temperature)


-- The Periodic Table: Organizing the Elements --– Nonmetals

• On the right side of the periodic table• Conduct electricity poorly or not at all and do not

have a luster• Exist as solids, liquids, and gases at room

temperature– Metalloids (which have properties of metals and

nonmetals) follow a zig-zag line• Boron, silicon, germanium, arsenic, antimony,

tellurium, polonium• Often used for computer chips and solar cells• Make up semiconductors because of their ability to

conduct electricity fairly but not extremely well

Chapter X: Nuclear Chemistry-- Exploring Radioactivity --

• Nuclide– General name for nucleus of atom– Neutrons are the only things that change in nuclides of

atoms of the same element• Types of radiation

– Alpha radiation ()• Alpha particles emitted by a radioactive source

– Positively charged particle emitted from certain radioactive nuclei

– Consists of two protons and two neutrons and is identical to the nucleus of a helium atom

Alpha particle


• Atomic numbers must add up, and mass numbers must add up. The new atomic number identifies the element as mercury (Hg).

• Properties of alpha radiation– Involves a helium nucleus– Charge: 2+ (no electrons to balance protons)– Mass: 4 amu– Penetrating power: 0.05 mm body tissue– Shielding: paper, clothing– Can be dangerous when ingested


• Examples

• Parent and daughter nuclides– Parent nuclide: initial nucleus– Daugher nuclide: resulting nucleus (aside from

emitted particle)» Less energetic than parent nuclide» Formed for stability

Parentnuclide

Daughternuclide


– Beta radiation ()• Beta particles emitted by a radioactive source

– Fast-moving electron emitted from certain radioactive nuclei

– Formed when a neutron decomposes into a proton and an electron

Beta particle

• Properties of beta radiation– Involves an electron– Charge: 1- (charge of an electron)– Mass: 1/1837 amu


– Penetrating power: 4 mm body tissue– Shielding: metal foil or thin pieces of wood– Process used in carbon-14 dating

• Examples


– Gamma radiation ()• Gamma rays emitted by a radioactive source

– High energy rays– Have no mass and no electrical charge

• Often emitted along with alpha or beta radiation by the nuclei of decaying radioactive atoms

• Properties of gamma radiation– Involves rays having no charge or mass– Charge: 0 (no charge in and of itself)– Mass: 0 amu

Gamma ray


– Penetrating power: through entire body– Shielding: lead, concrete (incomplete shielding)– Most dangerous of all radiation– Usually accompanied by alpha or beta particles– Similar in properties to X-rays

• Ionizing vs. nonionizing radiation– Ionizing radiation

• Has enough energy to change atoms and molecules into ions

• Examples– Alpha radiation– Beta radiation– Gamma radiation


– Nonionizing radiation• Does not have enough energy to change atoms and

molecules into ions• Examples

– Microwaves– Light waves– Radio waves

• Sources of radiation– The sun (cosmic radiation)– Food– Building materials– Background radiation (natural radiation)

• Carbon-14• Hydrogen-3 (tritium)


• Effects of ionizing radiation– Genetic mutations/annihilations

• Sensitivity of DNA• Carcinogenic consequences

– Threshold for radiation• Below: minimal damage• Above: radiation sickness

– Hair loss– Deformities– Cellular degradation/mutation


• Examples


• Examples

Chapter 6: Chemical Names and Formulas

• Introduction to Chemical Bonding• Representing Chemical Compounds• Ionic Charges• Ionic Compounds• Molecular Compounds and Acids• Summary of Naming and Formula Writing


Introduction to Chemical Bonding


-- Introduction to Chemical Bonding --Molecules and Molecular Bonding• Molecules

– Smallest electrically neutral unit of a substance that has all the properties of a substance

– Made up of two or more atoms that act as a unit– Examples

Water (H2O)

Methane (CH4)


-- Introduction to Chemical Bonding --• Molecular Compounds

– Compounds composed of more than one molecule– Caused by atoms combining to form compounds– Properties

• Relatively low melting and boiling points• Usually consist of two or more nonmetals• Examples

– Sugar (C12H22O11)– Methane (CH4)– Carbon dioxide (CO2)

• Noble gases are the only elements that exist in nature as isolated atoms; all others combine with atoms of different and same elements


-- Introduction to Chemical Bonding --Ions and Ionic Compounds• Ions

– Atoms or groups of atoms that have a positive or negative charge

– Formed by an electron being lost or gained by an atom– NOT caused by losing or gaining a proton– Example: Mg2+

• Formation of the positive charge comes from losing two electrons to become like a noble gas for stability

• 2+ because two electrons are lost


-- Introduction to Chemical Bonding --– Ions are usually formed to become as closed to a noble

gas and take the quickest route to get to being like a noble gas• Oxygen gains two electrons to be O2-

– (8 e- + 2 e- = 10 e-)– 10 e- matches with neon, a noble gas– Has a charge of 2- because it has 8 protons and

10 electrons, meaning 2 more negative charges– Gains two electrons instead of losing six

because it takes less energy


-- Introduction to Chemical Bonding --• Sodium loses one electron to be Na+

– (11 e- - 1 e- = 10 e-)– 10 e- matches with neon, a noble gas– Has a charge of 1+ because it has 11 protons

and 10 electrons, meaning 1 more positive charge

– Loses one electron instead of gaining seven because it takes less energy

– Names of ions• Cations

– Positively-charged ions or groups of atoms with overall positive charge

– Created by loss of electrons


-- Introduction to Chemical Bonding --– Named by taking the name of the element and

adding ion– Examples

» Na+: sodium ion» Mg2+: magnesium ion» Al3+: aluminum ion

• Anions– Negatively-charged ions or groups of atoms with

overall negative charge– Created by gain of electrons– Named by taking the first syllable of the name of

the element, adding –ide, and then ion


-- Introduction to Chemical Bonding --– Examples

» F–: fluoride ion» O2–: oxide ion » N3–: nitride ion

– Ionic compounds• Compounds composed of anions and cations• Properties

– Relatively high melting and boiling points– Consist of a cation and an anion– Usually consist of a metal cation and a nonmetal

anion– Examples

» Sodium chloride (NaCl)» Iron (III) oxide (Fe2O3)


Representing Chemical Compounds


-- Representing Chemical Compounds --Chemical formulas• Shows the number and type of atoms present in the smallest

representative unit of a substance• Types of chemical formulas

– Molecular formula• Chemical formula that shows the actual number and

kinds of atoms present in a molecule of a compound• Deals with molecular compounds, NOT ionic

compounds• Examples

– H2O: two atoms of hydrogen and one atom of oxygen

– CH4: one atom of carbon and four atoms of hydrogen


-- Representing Chemical Compounds --• Most elements, when not bonded with other

elements, exist alone as atoms.• Diatomic molecules

– Elements that do not exist as single atoms– Seven diatomic molecules

Seven Diatomic Molecules(KNOW THESE FOR THE REST OF THE TERM.)

H2 N2 O2 F2 Cl2 Br2 I2


-- Representing Chemical Compounds --– Formula unit

• Lowest whole-number ratio of ions in an ionic compound

• Deals with ionic compounds, NOT molecular compounds

• Examples– NaCl: one atom of sodium with one atom of

chlorine– Fe2O3: two atoms of iron with three atoms of

oxygen• Exist this way because of collections of patterns of

ions (see page 140)– Law of definite proportions

• In samples of any chemical compound, the masses of the elements are always in the same proportion.


-- Representing Chemical Compounds --• Ratios of compounds are always the same or they are

different compounds.– Law of Multiple Proportions

• Whenever two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers.


Ionic Charges


-- Ionic Charges --Monatomic ions• Ions consisting of only one atom• Examples

– Na+, Mg2+, Al3+

– N3–, O2–, F–

• Metallic atoms tend to lose electrons.– Metals have 1 to 3 valence electrons. To get to 8

electrons, it is easier for them to lose electrons, resulting in a positive (+) charge.

– Charges of metals• Group 1A: 1+ (lost one electron)• Group 2A: 2+ (lost two electrons)• Group 3A: 3+ (lost three electrons)

• Nonmetallic atoms tend to gain electrons.


-- Ionic Charges --– Nonmetals have 5 to 7 valence electrons. To get to 8

electrons, it is easier for them to gain electrons, resulting in a negative (–) charge.

– Charges of nonmetals• Group 5A: 3– (gained three electrons)• Group 6A: 2– (gained two electrons)• Group 7A: 1– (gained one electron)

• Some monatomic ions that have more than one charge– These ions are generally the transition metals.– The charges of ions formed by transition metals must be

indicated in parentheses using Roman numerals following the name of the metal.

– Examples• Cu+: copper (I) ion• Cu2+: copper (II) ion


-- Ionic Charges --– Exceptions

• Three transition metals ONLY have one charge, while three representative elements can have more than one charge.

• Transition metals with only one charge (and therefore no need to indicate charge in parentheses)– Ag+: silver ion– Cd2+: cadmium ion– Zn2+: zinc ion

• Representative elements that can have more than one charge (and therefore need charge indicated)– Lead (Pb)– Tin (Sn)– Antimony (Sb)


-- Ionic Charges --– Old naming scheme

• If an atom had only two main ions, the suffix -ous is given to the ion with the lesser charge and the suffix -icis given to the ion with the greater charge.

• Latin names are used instead of English generally.• Examples

– Cu+: cuprous ion (from cuprium)– Cu2+: cupric ion (from cuprium)– Fe2+: ferrous ion (from ferrum)– Fe3+: ferric ion (from ferrum)

• Rarely used anymore


-- Ionic Charges --Polyatomic ions• Ions consisting of more than one atom• Examples

– NH4+

– PO43-, SO4

2-, NO3-, Cr2O7

2-

• Refer to page 147 of your textbook for more polyatomic ions.• The suffix -ite is given to the polyatomic ion with one less

oxygen than another similar ion. The prefix per- indicates the most oxygen atoms and the prefix hypo- indicates the least oxygen atoms.– ClO4

- : perchlorate ion– ClO3

- : chlorate ion– ClO2

- : chlorite ion– ClO1

- : hypochlorite ion


-- Ionic Charges --• Other examples

– H2PO4–: dihydrogen phosphate

– NO2–: nitrite ion

– SiO32–: silicate ion

– NO3–: nitrate ion

– SO42–: sulfate ion

– PO43–: phosphate ion

– OH–: hydroxide ion– C2H3O2

–: acetate ion– NH4

+: ammonium ion– CO3

2–: carbonate ion


Ionic Compounds


-- Ionic Compounds --Binary compounds• Compounds consisting of two elements• For binary ionic compounds, remember that charges always

have to cancel out between cations and anions!!!!!!• Examples

– NaCl– Fe2O3

• Naming ionic binary compounds– First say the name of the cation and then the name of the

anion used to make the compound – Examples

• NaCl: sodium chloride– Na+ is the sodium ion– Cl- is the chloride ion


-- Ionic Compounds --• Fe2O3: iron (III) oxide

– Fe3+: iron (III) ion– O2-: oxide ion

– Naming entails knowing charges and balancing charges to write formulas.• "Knowing" example: iron (III) sulfide

– Iron has a charge of 3+ since it is iron (III).– Sulfide has a charge of 2- since sulfur gains two

electrons to become an ion.• "Balancing" example: iron (III) sulfide

– Since compounds must be neutral, one must balance the positive and negative charges.


-- Ionic Compounds --– Use least common multiples: since 6 is the LCM

between 2 and 3, 2 Fe3+ ions will make a charge of 6+ and three S2- ions will make a charge of 6-, balancing the charges out.

– Formula: Fe2S3– NOTE: Do not show charges in the formula since

they balance each other out


-- Ionic Compounds --Ternary ionic compounds• Ionic compounds consisting of three elements• Examples

– Ca(NO3)2– NH4NO3

• Same processes for naming and writing formulas as binary ionic compounds


Molecular Compounds and Acids


-- Molecular Compounds and Acids --Molecular compound naming and formula-writing• Molecular compounds are made up of nonmetals.• Examples

– H2O (water)– H2O2 (hydrogen peroxide)

• Because molecular compounds aren’t made up of ions and because the same elements can be combined in different ratios, molecular compounds are named using prefixes.

• Examples of naming– CO2: carbon dioxide– Cl2O8: dichlorine octoxide

• Rules for naming molecular compounds– If the first element listed has only one atom, no prefix is

needed.


-- Molecular Compounds and Acids --• Rules for naming molecular compounds

– First element• If the first element listed has only one atom, no prefix is

needed, so you just state the name of the element. (Example: NO2 would start off as “nitrogen” …)

• If the first element listed has more than one atom, add a prefix on the element name. (Example: N2O would start off as “dinitrogen” …)

– Second element• Add on a prefix (even if there is only one atom of that

element).• Keep the first syllable of the element and change the

ending to –ide. (Examples: “oxide,” “nitride,” etc.)• Prefixes


-- Molecular Compounds and Acids --– Examples of naming molecular compounds

• OF2: oxygen difluoride• SO3: sulfur trioxide• O2F2: dioxygen difluoride• N2O5: dinitrogen pentoxide• XeF4: xenon tetrafluoride

• Rules for writing molecular formulas– If you do not see a prefix before the element, there is only

one atom of that element.– If you see a prefix before the element, the prefix tells you

how many atoms there are of that element.– Examples of writing molecular formulas

• Carbon tetrachloride: CCl4• Phosphorus pentafluoride: PF5


-- Molecular Compounds and Acids --– 1: mono –– 2: di –– 3: tri –– 4: tetra –– 5: penta –– 6: hexa –– 7: hepta –– 8: octa –– 9: nona –– 10: deca –

• The vowel at the end of a prefix is often dropped when he name of the element begins with a vowel. (Example: O7 would be expressed as “heptoxide” instead of “heptaoxide”.


-- Molecular Compounds and Acids --Acids• Acids give off (dissociate) hydrogen ions (H+) in solution.• The stronger the acid, the more the ions dissociate in

solution.• Naming acids

– When the name of the anion ends in –ide, the acid name begins with the prefix hydro-. The –ide becomes an –ic. Then you add the word acid. (Example: HCl has the chloride ion, resulting in hydrochloric acid.)

– When the anion name ends in –ate, the –ate becomes an –ic. Then you add the word acid. (Example: HNO3 has the nitrate ion, resulting in nitric acid.)

• Formulas and names of common acids (KNOW THESE)– HF: hydrofluoric acid– HCl: hydrochloric acid


-- Molecular Compounds and Acids --– HBr: hydrobromic acid– HI: hydroiodic acid– HNO3: nitric acid– H2SO4: sulfuric acid– H3PO4: phosphoric acid– HC2H3O2: acetic acid– H2CO3: carbonic acid

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