chapter 3 atoms and the periodic table. reactivity groups reactivity metals as orbits are added the...

Chapter 3 Atoms and the Chapter 3 Atoms and the Periodic TablePeriodic Table

ReactivityReactivity

• Groups reactivity METALS

• As orbits are added the electrons move further away from nucleus and become easier to lose , thus as you move down the group the elements become more reactive

• Nonmetals need electrons therefore the elements closer to the nucleus are more capable of accepting electrons and the elements moving up the group become more reactive

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• Elements moving from left to right in a period becoming smaller due to the pull of the nucleus.

• Elements moving down a group become larger in size due to adding orbits

• The most reactive families have fewer electrons to lose or gain, therefore groups 1 and 17, why not 18?

• Group 2 more reactive than 3 ;;group 16 more reactive than 15

Gases on periodic tableGases on periodic table

• All of 18 are gases known as the noble gases

• Other gases that are found naturally

• H, N, O, F, Cl,

• When these gases appear by themselves they are written as diatomic gases

• What are Atoms?– defined - are tiny units that determine the

properties of all matter• an atom is the smallest part of an element that still

has the element's properties

– introduction• Democritus

– Greek philosopher• lived in the 4th century B.C.• suggest that the universe made of invisible units

called atoms• defined - Greek word meaning "unable to divide"• believed that the changes he observed was due to

the movement of the atoms• unable to provide the evidence needed to convince

people that atoms existed

– Atoms are the building blocks of molecules• John Dalton atomic theory in 1808

– English school teacher– widely supported due to supporting evidence– three parts

• every element is made of tiny unique particles called atoms that cannot be subdivided

• atoms of the same element are exactly alike• atoms of different elements can join to form

molecules

• What is in an atom?– introduction

• less than a 100 years after Dalton published his atomic theory scientist determined that atoms could be split further

• today we know there are many different parts of an atom but only three are used in everyday chemistry of most substances

– In the nucleus - dense center of the atom• protons - 1 positive charge with a mass of 1 amu

(atomic mass unit)• neutrons - 0 charge (neutral) with a mass of 1 amu

– Electron cloud - made of very tiny moving particles• electrons - 1 negative charge with very little mass 0

amu

• Atoms have no over all charge because they have an equal number of protons and electrons

– example He (helium) atom• 2 protons• 2 neutrons• 2 electrons• charge of 2 protons +2• charge of 2 neutrons 0• charge of 2 elections -2• total charge 0

• Models of the Atom– introduction

• like most scientific models and theories the model of the atom has been revised many time to explain each new discovery

– Bohr's model• Niels Bohr - Danish scientist in 1913

– electrons move in set paths around the nucleus like the planets orbit the sun

– each electron has a certain energy that is determined by its path around the nucleus

– energy level• the path of the possible energies an electron may have in

an atom– electrons must gain energy to move to a higher energy level

– Modern theory• by 1925 Bohr's model no longer explained all

observations– electrons no longer moved in definite paths– electrons behave like waves vibrating on a string than

like particles– impossible to determine the exact location, speed, and

direction• like a fan blade

– try to determine location by shading• the darker the shading the better the chance to find

an electron• the whole shaded region is called an electron cloud

• electrons are found in an orbital within each energy level (orbit or shell)

– orbital - the region in an atom where electrons are found• exist only when an electron occupies it

– four different kinds of orbitals– "s" orbital

• simplest• shaped like a sphere• only 1 orbital or orientation per orbit• can contain 2 electrons maximum

– "p" orbitals• dumbbell shaped• 3 different orbitals or orientations per orbit• x, y, and z axis• 2 electron in each orbital• 6 electrons maximum

– "d " orbitals• 5 possible orbitals or orientations per orbit• 2 electrons each orbital• 10 electrons maximum

– "f " orbitals• 7 possible orbitals or orientations per orbit• 2 electrons each orbital• 14 electrons maximum

• electrons start off occupying the lowest level then are added to the next highest energy level or orbit

– 1st energy level or orbit– contains only the "s" orbital– 2 electrons maximum

– 2nd energy level or orbit• contains

– "s" orbital• 2 electrons maximum

– "p" orbitals• 6 electrons maximum

– 8 electrons maximum for the energy level– 2 in the s and 6 in the p

– 3rd energy level or orbit• contains

– "s" orbital• 2 electrons maximum

– "p" orbitals• 6 electrons maximum

– "d " orbitals• 10 electron maximum

• 18 electrons maximum for the energy level– 2 in the s, 6 in the p, and 10 in the d

– 4th energy level• contains

– "s" orbital• 2 electrons maximum

– "p" orbitals• 6 electrons maximum

– "d " orbitals• 10 electrons maximum

– "f " orbitals• 14 electrons maximum

• 32 electrons maximum for the energy level– 2 in the s, 6 in the p, 10 in the d, and 14 in the f

– every atom has one or more valence electrons• valence electron - is an electron in the outer most

energy level of an atom• hydrogen has 1 valence electron (the least number)• neon has 8 valence electrons (the maximum

number)

• Page 76

• Questions 1-7

• Write questions and answers

3.2 A Guided Tour of the 3.2 A Guided Tour of the Periodic TablePeriodic Table

• Objectives– Relate the organization of the periodic table to the

arrangement of electron within an atom.– Explain why some atoms gain or lose electrons to

form ions.– Determine how many protons, neutrons, and

electrons an isotope has, given its symbol, atomic number, and mass number.

– Describe how the abundance of isotopes affects and element’s average atomic mass.

• Historical prospective– Developed by

• Dimitri Mendeleev– Russian chemist– in 1869– based on repeating properties and atomic mass– he arranged the known elements and left blank spaces

for unknown elements

– Henry Mosley• the first to group atoms by protons

• Organization of the Periodic Table– similar elements grouped together– makes it easier to predict the properties of an

element based on where it is in the periodic table

– elements represented by their symbols– order based on the number of protons the

atom has in its nucleus– Hydrogen has one proton and is the first

elements listed in the Periodic Table

– Period law• properties of elements tend to change in regular

pattern when elements are arranged in order of increasing atomic number, or number of protons in their nucleus

• atomic numbers equals the number of protons– increases from left to right and top to bottom

• Using the Periodic Table to determine electronic arrangement– Periods

• horizontal rows• 1-7• indicates the outer most energy level• Period 1

– 2 elements - H and He– has only 1 s orbital

• maximum of 2 electrons

• Period 2– starts with Li and ends with Ne– contains 1 s and 3 p orbitals

• Period 3– starts with Na and ends with Ar– contains 1 s and up to 3 p orbitals

• Periods 4 and 5– contains 1 s and up to 3 p and 5 d orbitals

• Periods 6 and 7– contains 1 s and up to 3 p, 5d, and 7 f orbitals

– Groups• vertical columns

– 1-18– have similar properties– have the same number of valence electrons

• Groups 1 and 2– electrons are going in the s orbital– Group 1

• H down to Fr• only one electron in the outer most energy level• 1 electron in the s orbital• these elements have 1 valence electron

– Group 2 • Be down to Ra• 2 electrons in the outer most energy level• 2 electrons in the s orbital and it is now full• these elements have 2 valence electrons

• Groups 13 to 18– are placing electrons in the p orbitals of the outer most

energy level– these elements have 3 to 8 valence electrons

• Group 13 elements have 3 valence electrons• Group 18 elements have 8 valence electrons• this is the maximum number of valence electrons• the p orbitals are full• Valence electrons equals the last digit of the Group

number

• Groups 3 to 12 – are placing electrons in the d orbital of the next lower

energy level– they have 2 valence electrons as far as this class is

concerned

• The Lanthanoid Series and Actinoid Series– are placing electrons in the f orbital of the energy level 2

place back– they have 2 valence electrons as far as this class is

concerned

• Atoms in Group 18 have full outer energy levels– 8 is the maximum for an outer level– except for level 1 He has 2– non reactive (inert)

• united video Elements of Chemistry: The Periodic Table (20:00 min.) http://www.unitedstreaming.com/search/assetDetail.cfm?guidAssetID=F59A819C-DB1B-48E6-90D7-DF3829C74230

• Elements are reactive because their outer most energy levels are only partially filled.

• Some Atoms Form Ions– Ionization

• defined - atoms that may gain or lose valence electrons so that they have a full outermost energy level

• no longer the same number of protons and electrons

• it has a net electrical charge

– ion • defined - an atom or group of atoms that has lost

or gained one or more electrons and therefore has a net electric charge

• cation– defined - an ion with a positive charge– example: Li has 1 valence electron

• 2 electrons in the 1st energy level• 1 electron in the 2nd energy level• when the valence electron is removed Li becomes a

positive ion Li+

• Li+ ion Li atom3 protons +3 3 protons +32 electrons -2 3 electrons -3 charge +1 0

• the other elements in Group 1 form +1 cations by having only one valence electron

• anion– defined - an ion with a negative charge– example: F has 7 valence electrons

• 2 electrons in the 1st energy level• 7 electrons in the 2nd energy level• easier to gain 1 electron than lose 7 electrons to

become a negative ion.• F- ion F atom

9 protons +9 9 protons +910 electrons -10 9 electrons -9charge - 1 0

• the other elements in Group 17 form -1 anions by having 7 valence electrons

• How Do the Structures of Atoms Differ– Atomic number

• defined - the number of protons in the nucleus of an atom

– remember atoms are always neutral because they have equal number of protons and electrons

• the simplest atom H has only 1 proton and 1 electron

– atomic number is 1

• the largest naturally occurring atom U has 92 protons and 92 electrons

– atomic number is 92

– Mass number• defined - the total number of protons and neutrons

in the nucleus of an atom• F has 9 protons and 10 neutrons for a mass

number (A) = 19• the mass number can vary from atom to atom of

the same element

– Isotopes• defined - atoms having the same number of

protons but different number neutrons

– example: H has 2 isotopes• the first is the protium

– the atom of H (the most common)– has only one proton and 0 neutrons– a mass number of 1

– the second isotope is Deuterium • sometimes called "heavy Hydrogen"• 1 proton and 1 neutron• a mass number of 2• only 1 out of every 6000 H are Deuterium

– the third isotope is Tritium• 1 proton and 2 neutrons• mass number of 3

– All three are hydrogen, only one proton, but have different masses due to the neutrons.

– Calculating the number of neutrons in an atom

• average atomic mass– defined - the weighted average of the masses of all

naturally occurring isotopes of an element

• This is found under the Symbol on the Periodic Table

– round this number to the nearest whole number– subtract the atomic number

– example: C• average atomic mass 12.011 = 12 mass number• atomic number - 6

number of neutrons 6• this is for the most common Carbon atoms (carbon -

12)• the isotopes for C will be those with different number

of neutrons like carbon - 14• Mass number 14

atomic number - 6neutrons 8

• Rules for Electron configuration– Find the total number of electrons (atomic

number).– Find the number of energy levels (the period

number).–

Draw the orbits

– Find the electrons in the last energy level.• For Groups 1& 2 use the Group number.• For Groups 13 - 18 use the last digit of the Group

number (3 - 8).• For He always 2 electrons.• For Group 3 - 12 assign 2 electrons

– Subtract the electrons from the total as you place them in their energy level.

– Fill in the inner energy levels with the remainder of the electrons starting with the first energy level.

• Use the following pattern when they are the inner energy levels.

– 1st energy level - 2 electrons– 2nd energy level - up to 8 electrons– 3rd energy level - 8 or 18 electrons– 4th energy level - 8, 18, or 32 electrons– 5th energy level - 8, 18, or 32 electrons– 6th energy level - 8 or 18 electrons

• Remember to subtract as you add them to their energy levels.

• Examples: Br K and Bi

• Page 85

• Questions 1-7

• Questions and answers

3.3 Families of Elements3.3 Families of Elements

• Objectives– Locate alkali metals, alkaline-earth metals,

and transition metals in the periodic table.– Locate semiconductors, halogens, and noble

gases in the periodic table.– Relate an element’s chemical properties to

the electron arrangement of its atoms.

• Groups are sometimes called families– each is unique yet share certain similarities– elements have common chemical and

physical properties• they have the same number of valence electrons

• How elements are classified• Metals

– the majority of all elements– most are

• solids• luster - shiny• ductile - can be stretched• malleable - can be shaped• good conductors of heat and electricity• form cations only

• Alkali Metals– the highly reactive metallic elements located

in Group 1 of the Periodic Table• Li, Na, K , Rb, Cs, and Francium• form cations with a +1 oxidation number• highly reactive• soft• luster - shiny

– Uses• NaOH (sodium hydroxide) used to manufacture

– paper– soap– synthetic fabrics– petroleum refining

• NaCl -table salt• KCl - table salt substitute• K - used in fertilizers• Na+ and K+ are important for proper functioning of

nerves in our bodies.

• Alkaline Earth Metals– reactive metallic elements located in Group 2

• Be, Mg, Ca, Sr, Ba, and Ra• form cations with a +2 oxidation number• less reactive than the Alkali Metals

– requires more energy to remove the 2nd electron than the 1st electron from an energy level

• light• good structural strength

– uses• Ca - animals shells, limestone, marble, bones, and

teeth• Mg - airplane frames, activates enzymes, flares,

Epson salt, and milk of magnesia

• Transition Metals or Elements– located in Groups 3 - 12– conducts heat and electricity like other metals– form multiple cations

• some up to 4 different cations

– frequently form colorful compounds such as rubies and emeralds

– uses• Ag – a better conductor than Au• Au does not corrode or tarnish under ordinary

conditions– great for connectors for computers and other electronic

devices

• Fe, Co, Cu and Mn play important roles in our body chemistry

• Hg the only metal that is a liquid at room temperature

– flows easily– does not stick to glass makes it good for thermometers

• Fe, Co, and Ni– in the same period– the only metals that can be magnetized

• Cu, Ag, and Au – in the same family– called the coinage metals

• radioactive isotopes– defined - nuclei are continually decaying to produce

different elements– used at times to detect cancer in the soft tissue of our

bodies

• Nonmetals– Except for H they are found on the right hand

side of the Periodic Table• some of the elements in Groups 13 - 16• all the elements in group 17 - 18• do not conduct electricity or heat• brittle• no luster - dull

– C• occurs in three forms naturally

– graphite– diamonds– fullerenes

• combine with other elements to form millions of compounds

– sugars, chlorophyll, gasoline, and rubber to name a few

– O, N, and S• common nonmetals• form anions of oxide-2, sulfide-2, and nitride-3

• most plentiful gases in the atmosphere are N and O

• S is an odorless yellow solid– many S compounds are known for their terrible smell

• rotten eggs, H2S

• skunk spray

• Halogens– Group 17– slightly reactive– forms salts with metals– form ions with a -1 oxidation number

– 4 of the 7 diatomic elements• 2 atoms per molecules as an element

• the other 3 are H2, N2, and O2

– Cl2 yellowish green poisonous gas used to kill bacteria in water

– F2 • most reactive non metal• a poisonous yellowish gas

– Br2 a dark red liquid

• Noble Gases– Group 18– non reactive gases exist as single atoms and

not as compounds– He lighter than air and used in balloons– Ne used in signs because of its reddish

orange glow– Ar used in light bulbs

• Metalloids or Semiconductors– have properties of metals and nonmetals– weak conductors of electricity and heat– solids– white or gray in color– B, Si, Ge, As, Sb, Te, and Po form a stair step

downward from left to right– metals are to the left of the metalloids and the

nonmetals are to the right

• Page 94

• Questions 1-7

• Questions and answers

3.4 Using Moles to Count Atoms3.4 Using Moles to Count Atoms

• Objectives– Explain the relationship between a mole of a

substance and Avogadro’s constant.– Find the molar mass of a n element by using

the periodic table.– Solve problems converting the amount of an

element in moles to its mass in grams, and vice versa.

• Counting Units– dozen (12 items)– bushel (32 qt container)– reams (500 sheets)– pairs (2)

• Mole is used for counting very small particles– abbreviated mol.– a collection of 602 213 670 000 000 000 000

000 particles– usually written as 6.022 x 1023 particles per

mole

– known as Avagodro's number or constant• named for Amedeo Avagodro

– an Italian that lived from 1776 - 1856– a lawyer interested in mathematics and physics– 1st to make a distinction between atoms and molecules

– This constant was determined by Joseph Loschmidt

• German physicist• in 1865

– 1 mole of popcorn kernels would cover the entire US to a height of 500 km (310 mi)

• not a good way to count popcorn

• Molar mass– defined - the mass in grams of 1 mol of a

substance– The molar mass of an element in grams is the

same as its average atomic mass in amu

– conversion factor• defined - a ratio equal to one that expresses the

same quantity in two different ways• 10 gumballs = 21.4 g• can be written as 10 gumballs / 21.4 g or 21.4 g /

10 gumballs• What is the mass of 50 gumballs?

– 50 gumballs x 21.4 g / 10 gumballs = 107 g

– p. 98 practice factors 1 - 3

– Relating moles to grams• 1 molar mass of element or 1 mol of element• 1 mol of element 1 molar mass of

element• Fe has 55.85 amu therefore 55.85 g Fe• 1 mol Fe

– determine the mass in grams of 5.5 mol of iron.• 5.50 mol Fe x 55.85 g Fe = 307 g Fe• 1 mol Fe

• p. 99 practice converting Amount to mass 1 & 4

– Converting mass to amount• Determine the amount of iron present in 352 g of

iron.• 352 g Fe x 1 mol Fe = 6.3 mol Fe• 1 55.85 g Fe

– How many moles are in 536 g of copper?• 536 g Cu x 1 mol Cu = 8.44 mol Cu• 1 63.55 g Cu

– How many moles are present in 12.1 g of sulfur.

• 12.1 g S x 1 mol S = .377 mol S• 1 32.07 g S

Page 100questions 1 – 9

• Write questions and answers (show work on the problems)