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Inside Atoms and Molecules

Chapter 3

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Are Atoms Really Unbreakable?• J.J. Thomson investigated a beam called a cathode ray• he determined that the ray was made of tiny negatively

charged particles we call electrons• his measurements led him to conclude that these

electrons were smaller than a hydrogen atom• if electrons are smaller than atoms, they must be pieces

of atoms• if atoms have pieces, they must be breakable• Thomson also found that atoms of different elements

all produced these same electrons

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The Electron

• Tiny, negatively charged particle

• Very light compared to mass of atom– 1/1836th the mass of a H atom

• Move very rapidly within the atom

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Rutherford’s Nuclear Model The atom contains a tiny dense center called the

nucleus– the volume is about 1/10 trillionth the volume of

the atom The nucleus is essentially the entire mass of the atom The nucleus is positively charged

– the amount of positive charge of the nucleus balances the negative charge of the electrons

The electrons move around in the empty space of the atom surrounding the nucleus

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(a) The results that the metal foil experiment would have yielded if the plum pudding model had been

correct; (b) Actual results

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A nuclear atom viewed in cross

section

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Structure of the Nucleus• The nucleus was found to be composed of two kinds

of particles• Some of these particles are called protons

– charge = +1

– mass is about the same as a hydrogen atom

• Since protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons

• The other particle is called a neutron– has no charge– has a mass slightly more than a proton

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Isotopes• All atoms of an element have the same number of protons• The number of protons in an atom of a given element is the

same as the atomic number– found on the Periodic Table

• Atoms of an element with different numbers of neutrons are called isotopes

• All isotopes of an element are chemically identical– undergo the exact same chemical reactions

• Isotopes of an element have different masses• Isotopes are identified by their mass numbers

– mass number = protons + neutrons

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Two isotopes of sodium

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The Modern Atom

• We know atoms are composed of three main pieces - protons, neutrons and electrons

• The nucleus contains protons and neutrons

• The nucleus is only about 10-13 cm in diameter

• The electrons move outside the nucleus with an average distance of about 10-8 cm– therefore the radius of the atom is about 105 times

larger than the radius of the nucleus

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(a) Continuous energy levels. (b) Discrete (quantized) energy levels.

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The difference between continuous and quantized energy levels

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Shell Structure for Electrons inAtoms• Electrons have a minimum energy, therefore they

do not crash into the nucleus

• The one electron of hydrogen is in an energy level that is closest to the nucleus

• The farther the energy level is from the nucleus, the higher its energy

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The shell model of the electrons in atoms

First Shell; can hold 2 electrons

Second Shell; can hold 8 electrons

Third Shell; can hold 18 electrons

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Bohr’s Model• Distances between energy levels decreases as the

energy increases– Electrons “orbit” the nucleus much like planets orbiting

the sun

• 1st energy level can hold 2e-1, the 2nd 8e-1, the 3rd 18e-1, etc.– farther from nucleus = more space = less repulsion

• The highest energy occupied shell is called the valence shell

• The other shells are called inner or core shells

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The relative sizes of the spherical first, second, and third electron shells

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The first four principal energy levels in the hydrogen atom

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The Modern Periodic Table

• Elements with similar chemical and physical properties are in the same column

• Columns are called Groups or Families

• Rows are called Periods

• Each period shows the pattern of properties repeated in the next period– Similar valence shell electron configurations

• Same numbers of valence electrons• Same orbital types• Different energy levels

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The periodic table

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Important Groups• Group 8 = Noble Gases

• He, Ne, Ar, Kr, Xe, Rn

• all colorless gases at room temperature

• very non-reactive, practically inert

• found in nature as a collection of separate atoms uncombined with other atoms

• Valence shells with 8

electrons lead to very

unreactive elements

• Helium is the only

Noble Gas with 2 electrons

in its valence shell

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Important Groups - Halogens

• Group 7A = Halogens• 7 e- in valence shell• very reactive

nonmetals—want to gain one electron to be like noble gases

• react with metals to form ionic compounds

• HX all acids

• Fluorine = F2 – pale yellow gas

• Chlorine = Cl2

– pale green gas

• Bromine = Br2

– brown liquid that has lots of brown vapor over it

– Only other liquid element at room conditions is the metal Hg

• Iodine = I2

– lustrous, purple solid

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Important Groups – Alkali Metals

• Group 1A• 1 e- in valence shell

• very reactive metals—want to lose one e- to be like noble gases

• react with nonmetals to form ionic compounds

• Li, Na, K are all violently reactive with water

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Covalent Bonds

• Typical of molecular substances

• Atoms bond together to form molecules– strong attraction

• Sharing pairs of electrons

• Molecules attracted to each other weakly

• Often found between nonmetal atoms

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The formation of a bond between two hydrogen atoms

H:H H—H

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Bond Polarity• Covalent bonding between unlike atoms

results in unequal sharing of the electrons– One end of the bond has larger electron density

than the other

• The result is bond polarity– The end with the larger electron density gets a

partial negative charge– The end that is electron deficient gets a partial

positive charge

H F••

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Probability representations of the electron sharing in HF

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Electronegativity• Measure of the ability of an atom to attract shared

electrons– Larger electronegativity means atom attracts more strongly– Values 0.7 to 4.0

• Increases across period (left to right) on Periodic Table

• Decreases down group (top to bottom) on Periodic Table

• Larger difference in electronegativities means more polar bond– negative end toward more electronegative atom

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Electronegativity values for selected elements

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Dipole Moment• Bond polarity results in an unequal electron

distribution, resulting in areas of partial positive and partial negative charge

• Any molecule that has a center of positive charge and a center of negative charge in different points is said to have a dipole moment

• The dipole moment effects the attractive forces between molecules and therefore the physical properties of the substance

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The three possible types of bonds: (a) a covalent bond; (b) a polar covalent bond; and (c) an ionic bond

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(a) The charge distribution in the water molecule. (b) The water molecule behaves as if it had a

positive end and a negative end

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Polar water molecules are strongly attracted to each other

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Lewis Symbols of Atoms and Ions• Also known as electron dot symbols• Use symbol of element to represent nucleus and inner

electrons• Use dots around the symbol to represent valence electrons

– put one electron on each side first, then pair

• Elements in the same group have the same Lewis symbol– Because they have the same number of valence electrons

• Cations have Lewis symbols without valence electrons• Anions have Lewis symbols with 8 valence electrons

Li• Be• •B• •C• •N• •O: :F: :Ne:• •

•

• • • •

•• •• •• ••

••

Li• Li+1 :F: [:F:]-1

•

•• ••

••

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Writing Lewis Structures of Molecules • Count the total number of valence electrons from all the

atoms• Attach the atoms together with one pair of electrons

– A line is often used as shorthand for a pair of electrons that attach atoms together

• Arrange the remaining electrons in pairs so that all hydrogen atoms have 2 electrons (1 bond) and other atoms have 8 electrons (combination of bonding and nonbonding)

• Occasionally atoms may violate this rule– Nonbonding pairs of electrons are also know as Lone Pairs

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Electrical Nature of Matter• Most common pure substances are very poor conductors

of electricity– with the exception of metals and graphite

– Water is a very poor electrical conductor

• Some substances dissolve in water to form a solution that conducts well - these are called electrolytes

• When dissolved in water, electrolyte compounds break up into component ions– ions are atoms or groups of atoms that have an electrical charge

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(a) Pure water does not conduct a current; (b) Water containing a dissolved salt conducts electricity

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Polar water molecules interact with the positive and negative ions of a salt

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Ions• ions that have a positive charge are called cations

– form when an atom loses electrons

• ions that have a negative charge are called anions – form when an atom gains electrons

• ions with opposite charges attract – therefore cations and anions attract each other

• moving ions conduct electricity• compound must have no total charge, therefore we

must balance the numbers of cations and anions in a compound to get 0 total charge

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Atomic Structures of Ions• Metals form cations• For each positive charge the ion has 1 less electron than the neutral

atom– Na = 11 e-, Na+ = 10 e-

– Ca = 20 e-, Ca+2 = 18 e-

• Cations are named the same as the metalsodium Na Na+ + 1e- sodium ioncalcium Ca Ca+2 + 2e- calcium ion

• The charge on a cation can be determined from the Group number on the Periodic Table for Groups IA, IIA, IIIA– Group 1A +1, Group 2A +2, (Al, Ga, In) +3

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Atomic Structures of Ions• Nonmetals form anions• For each negative charge the ion has 1 more electron

than the neutral atom– F = 9 e-, F- = 10 e-

– P = 15 e-, P3- = 18 e-

• Anions are named by changing the ending of the name to -ide

fluorine F + 1e- F- fluoride ionoxygen O + 2e- O2- oxide ion

• The charge on an anion can be determined from the Group number on the Periodic Table– Group 7A -1, Group 6A -2

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Electron Arrangements And Ion Charge• We know

– Group 1A metals form ions with +1 charge– Group 2A metals form ions with +2 charge– Group 7A nonmetals form ions with -1 charge– Group 6A nonmetals form ions with -2 charge– Group 8A nonmetals do not form ions, in fact

they are extremely unreactive

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Common Ions with Noble Gas Configurations in Ionic Compounds

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• Representative metals lose their valence electrons to form cations

• Nonmetals gain electrons so their valence shell has the same electron arrangement as the next noble gas

• There have to be enough electrons from the metals atoms to supply the needed electrons for the nonmetal atoms– Allows us to predict the formulas of ionic compounds

• In Polyatomic ions, the atoms in the ion are connected with covalent bonds. The ions are attracted to oppositely charged ions to form an ionic compound

Electron Arrangements and Ionic Bonding

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Ionic Bonds• Results from reaction between Metal and Nonmetal• Metal loses electrons to form cation, Nonmetal gains electrons to

form anion• Ionic bond is the attraction between a positive ion and negative ion• Larger Charge = Stronger Attraction• Smaller Ion = Stronger Attraction• No bond is 100% ionic!!• Electrostatic attraction nondirectional

– no direct anion-cation pair, No ionic molecule• chemical formula is empirical formula, simply giving the ratio of ions based on

charge balance• Ions arranged in a pattern called a crystal lattice

• maximizes attractions between + and - ions

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Bonding and Structure of Ionic Compounds• Crystal Lattice = geometric pattern determined by the size and

charge of the ions• Anions larger than cation

– Almost always– Anions larger than parent atom, Cations smaller than parent atom

• Anions generally considered “hard” spheres packed as efficiently as possible, with the cations occupying the “holes” in the packing

• Arrangement results in each cation being surrounded by as many anions as will fit– And visa versa– Maximizes attractions between ions

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The structure of lithium fluoride

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Relative sizes of some ions and their parent atoms