Download Chapter 6 the periodic table

Chapter 6 the periodic table
Organizing the elements
 Early chemists grouped elements according
to their properties
 J.W. Dobereiner
– Grouped elements into triads (groups of three)
– Had similar properties
– Not all elements could be put into triads
Mendeleev’s Periodic table
 Dmitri Mendeleev
– Developed his own periodic table of elements.
– Arranged elements into a periodic table with
groups based on repeating properties.
– Arranged the element in his periodic table in
order of increasing atomic mass.
Modern periodic table
 Henry Moseley
– Determined the atomic number for each of the
known elements
– Since atomic number is unique to each
elements, increasing atomic numbers are used
to arrange the modern periodic table.
 Arranging the elements into periods (rows)
give rise to Periodic Law – when elements
are arranged in order of increasing atomic
number, there is a periodic repetition of their
physical and chemical properties.
Metals, nonmetals, metalloids
 Three large groups that elements can be
placed into.
 As you move from left to right in a period
elements become less metallic and more
nonmetalic
Metals
 Good conductors of heat and electrical current
 Cleaned or freshly cut metals have high luster or
sheen. Reflect light
 Many are ductile (can be pulled into wire)
 Most are malleable – can be hammered into a
sheet
 All are solids at room temp. except mercury
 About 80% of all elements
Nonmetals
 Upper right hand side of periodic table to the right
of the step
 Most nonmetals are gases at room temperature.
Some are solids and one is a liquid
 Hard to generalize because there are many
different properties
 Opposite of metal
– Poor conductors
– Solid nonmetals tend to be brittle.
Metalloids
 Border bold stair line
 Properties similar to both metals and
nonmetals
 Properties depend on the situation the
element is in.
 Change conditions that the element is in
Squares in the periodic table
 Periodic table displays the symbols and
names of elements along with information
about the structure of the element
Groups in the periodic table
 Group 1A – Alkali metals
 Group 2A - Alkaline earth metals
 Group 7A – Halogens
– Can be prepared from their salts
 Electrons in elements
– # of electrons in the outer energy level
determines the properties of elements
Noble gases
 Group 8A
 Inert gases, rarely take part in reactions
 Filled s and p orbitals for their outer energy
level
Representative elements
 1A through 7A
 Wide range of physical and chemical
properties
 s and p orbitals of highest energy level are
part filled
 Its group number equals the # of electrons
in the highest occupied energy level
Transition elements
 Those elements found in the B groups
 Separate group A on left from group A
elements on the right
 Transition metals have electrons in the d
orbitals
 Inner transition metals have electrons in the
f orbitals
– Used to be called rare earth metals
Periodic trends
 Atomic size
– Look at molecules of identical atoms
– Determine the space between nuclei
– Use this determine the atomic radius
 One half the distance between the nuclei of two atoms of the
same element when the atoms are in a gaseous state
 Expressed in picometers ( one trillionth of a meter)
– In general atomic size increases from top to bottom
within a group and decrease from left to right across a
period
Ions
 Atom or group of atoms that have a positive or
negative charge.
 Atoms are neutral
 Ions form when electrons are transferred between
atoms
 Metals tend to form ions by losing electrons and
forming a positive ion.
 An ion with a positive charge is called a cation.
 Nonmetals tend to form negative ions by gaining
one or more electrons
 An ion with a negative charge is called an anion
Trends in ionization energy
 Energy required to remove an electron from
an atom
– Measured with the element in gaseous form.
– 1st ionization energy – energy needed to
remove first electron
– Tends to decrease from top to bottom within a
group, and increase from left to right across a
period.
Trends in ionic size
 Metals tend to lose electrons to form a
cation
 Nonmetals tend to gain electrons and form
an anion
– Cations are always smaller than the atoms from
which they form. Anions are always larger than
the atom from which they form.
Trends in electronegativity
 Ability of an atom of an element to attract electron
when the atom is in a compound.
 In general, electronegativity values decrease from
top to bottom within a group. For the
representative elements, the values tend to
increase from left to right across a period. Noble
gases not included.
 Metals to the left have low values
 Nonmetals on the right have high values
(excluding noble gases)
 transition metals do not exhibit a trend.
Chapter 7
Ions
Valence electrons
 Mendeleev – groups based off of properties
 Properties determined by # of valence
electrons
– Electrons in the highest most energy level
– Equal to the elements group number
– Helium exception
Electron dot structures
 AKA Lewis dot structures
 Diagrams to show the valence electrons
– Only include the valence electrons because
they are the only one that take part in chemical
bonds. See pg 188 table 7.1
Octet rule
 Noble gas – unreactive – Stable
 Other element try to have the same electron
configuration.
– Form compound to achieve this configuration.
 Octet rule because Noble gases have eight
valence electrons.
Octet rule
– Metals lose valence electrons – complete octet
in the next lowest energy level
– Nonmetals gain electrons or share electrons w/
other nonmetals to complete octet
Formation of cations
 When metals form a cation the name is the
same as the atom
– Potassium atom ------- potassium ion
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Properties are not the same however
Group 1A elements form 1+ cations
Group 2A elements form 2+ cations
The charges of cations from the transition
metals vary
Formation of Cations Continued
 Within the transition metals there are
exceptions to the octet rule see pg 190
– Ag
– Make pseudo noble gas configuration 18
electrons but all of the outer orbitals are filled so
it is relatively favorable.
Formation of anions
 Atom or group of atoms with a negative
charge
– Gained 1 or more electrons
– The name of an anion is different then the atom
it formed from
 Chlorine forms a chloride ion
 Oxygen forms oxide
Ionic bonds and Ionic compounds
• Ionic compound – made of anions and
cations.
• Neutral ( +1 -1 = 0)
• Held together with an ionic bond
• Attraction of opposing charges.
 Chemical formula- kind and # of atoms in
the smallest representative unit of a
substance
– In ionic compounds we look at the formula unit
 Lowest whole number ratio of ions in an ionic
compound
 For covalent bonds use the molecular formula
Properties of ionic compounds
 Most are crystal structures at room temp.
 Generally have high melting points
 Conduct electricity when melted or dissolved in
water.
 As crystals ionic compounds have a coordinate
number
– # of ions of opposite charge that surround the ion in the
crystal
– Determines the shape of the crystal
Metallic bonds
 Valence electrons in metals are in a sea of
electrons
– Electrons are mobile and can drift from area to
area within the metal
– Electrons surround the metal cation. (not a
neutral atom)
 Metallic bonds – attraction between
electrons and metal cations that hold metals
together
Sea of electrons
 Sea of electrons explains:
– Conductivity – Free floating electrons
 As one electron enters one leaves.
– Malleability/ductility – sea of electrons keep
metal cations away from one another
 In ionic compounds like charges get close together
and repel –fracturing crystal.
Crystalline structure
 In pure form metals form very simple
crystals
 Three examples
– Body centered ( 8 neighbors)
 Na, K, Fe, Cr, W
– Face centered (12 neighbors)
 Cu, Ag, Au, Al, Pb
– Hexagonal close-packed (12 neighbors)
 Mg, Zn, Cd
Alloys
 Mixture of one or more elements with at
least one being a metal.
– Brass – mixture of copper and zinc
 Superior to the parent materials
– Steel – corrosion resistance, ductility, hardness,
and toughness
Types of alloys
 Depending on the size of the atom
determines how they fit into the crystal
– If the atoms are close to the same size they fit
into the crystal - substitutional alloy
 Brass
– Different sizes – smaller atom fit into the spaces
between larger atoms – interstitial alloy
 Steel
Naming Ions
 Monatomic ions
– Single atom with a positive or negative charge
resulting from the lose or gain of one or more
valence electrons.
Cations
 Metallic elements tend to lose valence
electrons
– 1A  1+ ions (Li, Na, K, Rb, Cs, Fr)
– 2A  2+ ions (Be, Mg, Ca, Sr, Ba, Ra)
– 3A  3+ ions (Al the only common one)
– The charges of group A metal ions is equal to
the group #
– Name stays the same.
Anions
 Nonmetals tend to gain electrons and
become negatively charged
 Charge can be determined by subtracting 8
from the group number
– Cl group 7A 7-8= -1 charge
 Name changes – ends in –ide
 Groups 4A and 8A usually don’t form ions.
Ions from transition metals
 Many transition metals form more than one
cation with different charges.
 Iron forms Fe2+ (lose of two electrons) and
Fe3+ ( lose of three electrons)
– See table 9.2 on page 255
 Two naming methods
– Stock method
– Classical method
Stock method of naming
 Roman numeral is placed in parenthesis
after the name of the element to indicate the
value of the charge.
– Tin(IV)
 there is no space between name & number
Classical method of naming
 Root of Latin name is used with a suffix
– -ous ending used with lower charge
– -ic ending used with higher charge
– Ferrum  Ferrous and Ferric
 Look at names to figure out the symbol (Ferrous)
 These names do not actually tell you the charge.
Polyatomic ions
 Ions composed of more than one atom.
– Structures on page 257
 The names of most end in -ite or –ate.
– If they do the ion contains oxygen
– -ite ending for lower # of oxygen in the ion
– -ate ending for higher # of oxygen in ion.