CHAPTER 6 PERIODIC TABLE
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CHAPTER 6 PERIODIC TABLE
• Chemists used properties of elements to sort them into groups.
• early classifications grouped elements into “triad” with similar properties.
• ie. Cl, Br, I – similar properties
▪ reacted easily with metals
▪ had values (amus) that grouped them.
Mendeleev’s Periodic Table
• 1869 published a table of elements
• Meyer published nearly identical table
• Mendeleev – elements are arranged into groups based on a set of repeating properties. He arranged elements in order of increasing atomic mass.
• 6 years after his death Mosely discovered concept of atomic number – better method of organizing elements.
• Modern Periodic Table:
▪ vertical columns are called groups or families.
▪ elements of the same groups have similar chemical and physical properties.
▪ Elements of the same group have similar electron configurations.
▪ all members of a group have the same valence configuration but different principal quantum numbers.
▪ the number of valence electrons = the Group number.
▪ hydrogen occupies a unique position – it does not fit naturally into any Group.
▪ Horizontal rows are called periods.
▪ elements of the same period do not have much in common EXCEPT that the energies of their outermost electrons are similar.
▪ electrons are added one at a time moving from left to right across a period.
• 7 rows or periods:
• each period # corresponds to a principal energy level in which valence electrons are located.
o potassium is in period 4 so it has valence electrons in the 4th PEL.
• period 1 – 2 elements valence electrons in PEL 1
• period 2 – 8 elements valence electrons in PEL 2
• period 3 –
• period 4 – 18 elements valence electrons in PEL 3
• period 5
• period 6 – 32 elements valence electrons in PEL4
• properties within a period change as you move left to right.
• pattern of properties within a period repeats as you move from one period to the next.
• number of valence electrons increases from left to right
• each period ends with a noble gas.
• groups 1 and 2 most active
• metallic properties increase from the top to the bottom.
▪ Properties of Metals:
• solid at room temperature – except mercury
• have densities greater than water (alkali metals- group 1) will float.
• malleable – can be hammered into shape
• ductile – drawn or pulled into a wire
• luster – shiny
• good conductors of heat and electricity because their valence electrons are very “mobile”
• low ionization energy and low electronegativity
• lose electrons to form positive ions with smaller radii
• Standard Properties:
▪ many are gases network solids at room temperature
▪ NOT malleable or ductile – tend to be brittle in solid phase.
▪ LACK luster – dull
▪ high ionization energy and high electronegativity
▪ poor conductors of heat and electricity
▪ tend to gain electrons to become negative ions with radii larger than their atoms.
• B,Si,Ge,As,Sb,Te, and At
• adjacent to diagonal staircase
• intermediate type of element
• display both properties of metals and nonmetals.
PROPERTIES OF ELEMENTS:
• Ionization energy
• Atomic Radii
• Ionization Energy:
▪ the amount of energy needed to remove the most loosely bound electron from an atom
▪ measured when atom is in gaseous state.
• atoms with more than one electron have more than one ionization energy.
• removing electrons produces cations = positively charged ions.
Some compounds are composed of particles called ions.
Ion is an atom or group of atoms with a positive or negative charge.
Positive and negative ions form when electrons are transferred between atoms.
cation = ion with a positive charge
anion = ion with a negative charge.
TRENDS IN IONIZATION ENERGY
o periods - values from left to right increase – increase in the number of protons so more energy is needed to remove them from the atom. Hence, increasing ionization energy. You need more energy to remove electrons.
o Group – top to bottom ionization energy decreases. Valence electrons are at higher energy levels, farther from the nucleus, easier to remove them.
• atomic radius is one half of the distance between the nuclei of two atoms of the same element when the atoms are joined or half the distance between two adjacent atoms in a crystal.
TRENDS IN ATOMIC RADII:
• period - left to right – decrease in atomic radii
▪ metals have larger radii than non metals
▪ valence electrons of a period are the same energy level but the number of protons increases attracting them and the radii decreases.
• group – top to bottom – atomic size increases.
▪ more levels reduces the attractive force of the nucleus
IONIC RADII: distance from the nucleus to the outer energy level of the ion.
• metals - lose valence electrons and become positive ions
• nonmetals tend to gain electrons and become negative ions
• compare size of radii before and after losing or gaining electrons.
• the ability of an atom of an element to attract electrons when the atom is in a compound.
• groups - electronegativity values decrease from top to bottom. so highest value at the top of group. Attraction for bonded electrons is less toward the bottom of the group.
• period – electronegativity increases from left to right. Metals tend to have low values ; non metals higher.
• values can be used to predict type of bond that can be formed
▪ cesium least electronegative
• least tendency to attract electrons
• tends to react and lose electrons and form + ions.
▪ fluorine is most electronegative
• strong tendency to attract electrons
• forms a negative ion
• some elements can be found uncombined in nature – as an atom.
• oxygen = O2
• noble gases – free elements
• others so reactive – never uncombined
• groups 1,2 and 17.
Group 1 – Alkali Metals
• 1 valence electron
• form ions with +1 charge
• reactivity of elements increase as you go down a group.
• most reactive metals
• francium is most reactive metal
Group 2 Alkaline earth metals
• easily lose electrons
• never found in nature in atomic state (compounds)
• 2 valence electrons
• form +2 ions
• less reactive than alkali metals
• low ionization energies
• low electronegativity values
• lose electrons to form ionic bonds
• top to bottom – reactivity increases
• groups 3-12
• outermost d orbitals are being filled
• hard solids
• high melting points – except mercury
• multiple oxidation states
• far less reactive than metals in group 1 and 2
• often form ions that have color
• alloys – 2 or more metals blended together.
Inner Transition Metals:
• Lanthanides – very reactive can’t use commercially
• Actinides – reactive
• metals and metalloids (semi metals)
• 3 valence electrons
• forms +3 ions
• form covalent bonds
• do not form ions
• change from nonmetallic to metallic from top to bottom
• metals/semimetals/ nonmetals
• 5 valence electrons
• form -3 charge ions
• 6 valence electrons
• -2 charged ions
• metals/ nonmetals/ metalloids(semimetals)
Group 17 – The Halogens
• F2, CL2 – gases
• Br2 – liquid
• I2 – solid
• metals/nonmetal/ metalloid
• usually in the form of salts
Group 18 – Noble Gases
• nonmetals only
• inert gases
• contain 8 valence electrons
• not reactive @STP
• Kr and Xe can be reactive @ extreme conditions
• monoatomic gases
• found in group I because it has 1 valence electron
• nonmetal at STP in gas form
• colorless and odorless
• bonds with carbon to form organic compounds
• forms ions with a +1 or -1 charge.
Elements in periodic table lie in “blocks”
• elements generally in same block have same characteristics.
• s-block = all metals- groups 1A and 2A and helium
• p-block = some metals and non-metals – metalloids. – groups 3A,4A,5A,6A,7A,8A except helium.
o metals and metalloids are solid with the exception of mercury
o nonmetals include solids, a liquid (bromine) and gases.
• d – block = transition metals -
• f – block = inner transition metals
• period = principal energy level
• period 1 and 2 - s and p sublevels are filled
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