Chapter 6 The Periodic Table 682mf

Chapter 6: The Periodic Table

Lesson 1 Organizing the Elements

Forerunners ● Only 13 elements identified by 1700 o Scientists suspected existence of others, but were unable to isolate elements from their compounds ● Metals (Cu, Ag, Au) since prehistoric times ● Identified hydrogen, oxygen, nitrogen, carbon between 1765 and 1775

Forerunners ● Atomic spectroscopy introduced in the 1800s ● Elements identified by unique line spectra o Light emitted when excited electron returns to its ground state ● Scientists doubled number of known elements

Forerunners: Law of Triads Johann Wolfgang (JW) Dobereiner (1780-1849): ● German chemist that introduced Law of Triads

Forerunners: Law of Triads ● Classified elements in groups of 3 ● Properties of middle element were approximate averages of the properties of the first and third element ● Not all known elements could be organized into triads

Forerunners: Law of Triads Chlorine, Bromine, and Iodine formed a triad.

Average of mass of Cl and mass of I = (35.5 + 126.9)/2 = 81.2 amu Bromine has an atomic mass of 79.9 amu!

Forerunners: Law of Octaves J.A.R Newlands: ● Introduced Law of Octaves ● 62 known elements so far ● Elements arranged by increasing atomic mass exhibited periodic pattern every 8 elements o First and eighth elements had similar properties, second and ninth, third and tenth, etc.

Forerunners: Law of Octaves ● Based name on 8 notes of musical scale so colleagues did not take Newlands seriously! ● It took 20 years to get credit for his idea and 100 years for scientists to truly accept it

The Periodic Table: Mendeleev vs. Meyer Dmitri Mendeleev and Lothar Meyer published nearly identical classifications of elements in the late 1800s.

The Periodic Table: Mendeleev vs. Meyer Mendeleev published first so he received credit. ● Wanted to help students learn elements (more than 60 at the time) more easily ● Wrote elements’ names and properties on cards and arranged them in various ways ● Developed before scientists knew about the structure of atoms

The Periodic Table: Mendeleev vs. Meyer ● Increasing atomic mass: o Periodic repetition of properties ● Predicted existence of 3 new elements o Ekasilicon was one of these -- 15 years later, scientists discovered it and named it Germanium o Lead to acceptance of his arrangement

Atomic Number Henry (HGJ) Moseley: ● Post-doc in Rutherford’s lab ● Some elements in Mendeleev’s table were still out of order

Atomic Number ● 1913: Moseley observed that metals produce different frequencies of x-rays when struck with energetic electrons o Hypothesized that this arose from different fundamental properties (amount of positive charge in the nucleus) ● Determined the correct way to organize the periodic table was by atomic number, not atomic mass

The Periodic Law ● Elements within a column (group) have similar properties ● Properties within a row (period) change as you move from left to right o Pattern of properties within a period repeats as you move from one period to the next

The Periodic Law When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern.

Metals, Nonmetals, Metalloids ● Elements can be grouped into 3 broad classes based on their general properties o Metals o Nonmetals o Metalloids (Semimetals) ● Across a period, properties of elements become less metallic and more nonmetallic

Metals, Nonmetals, Metalloids Metals: ● About 80% of elements are metals ● Good conductors of heat and electric current ● Freshly cleaned or cut metal surface has a luster or sheen ● All solids at room temperature (exception: mercury)

Metals, Nonmetals, Metalloids Metals: ● Ductile o Can be drawn into wires ● Malleable o Can be hammered into thin sheets

Metals, Nonmetals, Metalloids Nonmetals: ● Greater variation in physical properties ● Most are gases at room temperature, a few are solids, one is a liquid (bromine) ● Generally poor conductors of heat and electric current o Exception: carbon ● Solid nonmetals tend to be brittle

Metals, Nonmetals, Metalloids Metalloids: ● Properties similar to both metals and nonmetals ● Behavior can often be controlled by changing conditions o Example: pure silicon is a poor conductor of electric current but if mixed with a small amount of boron, it becomes a good conductor of electric current

Example Which of the following sets of elements have similar physical and chemical properties? Explain how you determined your answer. a) oxygen, nitrogen, carbon, boron b) strontium, magnesium, calcium, beryllium c) nitrogen, neon, nickel, niobium

Example Identify each element as metal, non-metal, or metalloid.

1) 2) 3) 4)

gold silicon sulfur barium

Example What does atomic number tell you about atoms of an element? Why is is better to use atomic number than atomic mass for organizing the periodic table?

Lesson 2 Classifying the Elements

Squares in the Periodic Table ● The periodic table displays the symbols and names of the elements and information about their structure

Atomic number Element symbol Element name Atomic mass

Electron Configurations in Groups ● Role of e-s is key in determining properties of elements ● Elements can be sorted into the following based on their e- configuration: o Noble gases o Representative elements o Transition metals o Inner transition metals

Group 8A: Noble Gases ● Occur in trace amounts in atmosphere ● Sometimes called inert gases o Very unreactive o 1962: Neil Bartlett made Xenon Tetrafluoride ● s and p sublevels are completely filled with e-s

Groups 1A-7A: Representative Elements ● Display a wide range of physical and chemical properties o Metals, nonmetals, metalloids o Many solids, a few gases, one liquid (bromine) at room temperature

Groups 1A-7A: Representative Elements ● s and p sublevels of the highest occupied energy level are not filled o Group 1A: 1 e- in highest occupied level o Group 4A: 4 e-s in highest occupied level ● Group number always equals number of e-s in highest occupied energy level

Group B: Transition Elements ● Separate group A elements o Transition metals and inner transition metals o Classified based on e- configuration

Group B: Transition Elements Transition metals: ● Displayed in main body of periodic table ● Highest occupied s sublevel and nearby d sublevel contain e-s ● Characterized by the presence of d orbitals ● Include copper, silver, gold, and iron

Group B: Transition Elements

Inner transition metals: ● Appear below main body of periodic table ● Highest occupied s sublevel and a nearby f sublevel generally contain e-s ● Used to be referred to as rare earth elements o Misleading because some are more abundant than other elements

Blocks of Elements ● Consider both e- configuration and position of elements on periodic table o Another pattern emerges

s-block Elements ● ● ● ●

Valence e-s only occupy s-orbitals Group 1A: 1 valence e- in an s-orbital Group 2A: 2 valence e-s in an s-orbital s-block contains only 2 groups (plus helium) because sorbitals can only hold 2 e-s

p-block Elements ● Valence e-s fill p-orbitals ● Group 3A to 8A (except helium) ● First energy level has no p-sublevels so the first period (row) has no pblock elements ● p-orbitals begin with 2p so p-block begins in 2nd period of the table (with Boron) ● p-block is 6 elements wide because p-orbitals hold maximum of 6 e-s

d-block Elements ● 10 elements wide because d-orbitals holds up to 10 e-s o Transition metals ● First d-orbital (3d) begins filling with Scandium ● Electrons are added to a d sublevel with a principal energy level that is 1 less than the period number

f-block Elements ● 4f orbital begins filling in the 6th period of the table ● 14 elements wide because f-orbitals hold a maximum of 14 e-s o Inner transition metals ● Electrons don’t fill f-orbitals in a regular pattern! They do not progress sequentially like other orbitals! ● Principal energy level of the f sublevel is 2 less than the period number

Example Use the diagram from the last slide to write configurations for: a) Carbon

b) Strontium c) Vanadium

Example List the symbols for all the elements whose econfiguration end as follows. Each n represents an energy level. a) ns2 b) ns2np5 c) ns2np6nd2(n+1)s2

Lesson 3 Periodic Trends

Trends in Atomic Size Atomic radius: ● One half the distance between the nuclei of 2 atoms of the same element when the atoms are ed ● Measured in picometers (1012 pm = 1 m)

Trends in Atomic Size Atomic size tends to: ● Increase from top to bottom within a group ● Decrease from left to right across a period

Example Why do you think atomic size increases from top to bottom, but decreases from left to right?

Group Trends in Atomic Size ● For alkali metals and noble gases, atomic radius increases as atomic number increases ● As atomic number increases within a group: o The charge on the nucleus increases  Increases in positive charge draws e-s closer to nucleus o The number of occupied energy levels increases  Shields e-s in the highest occupied energy level from the attraction of protons in the nucleus

Group Trends in Atomic Size ● Shielding effect is greater than the effect of the increase in nuclear charge o Atomic size increases

Periodic Trends in Atomic Size ● In general, atomic size decreases from left to right across a period o Each element has 1 more p+ and 1 more e- than preceding element o Across a period, e-s added to same principal energy level

Periodic Trends in Atomic Size ● Shielding effect is constant for all elements in a period o Increasing nuclear charge pulls e-s in highest occupied energy level closer to nucleus o Atomic size decreases

Example Explain why fluorine has a smaller atomic radius than both oxygen and chlorine.

Ions ● Many periodic trends related to how elements behave during chemical reactions

Ions ● Ion: o An atom that has a positive or negative charge o Forms when e-s are transferred between atoms

Ions Cation: Metallic elements tend to form positive ions by losing 1 or more e-s from highest energy level

Anion: Nonmetallic elements tend to form negative ions by gaining 1 or more e-s

Trends in Ionization Energy Ionization energy: ● Energy required to remove an e- from an atom ● Measured when element is in its gaseous state ● Helps predict what ions elements will form

Trends in Ionization Energy ● Energy required to remove first e- is the first ionization energy o Tends to decrease from top to bottom in a group o Tends to increase from left to right across a period

First Ionization Energies

Ionization Energies (for Z=1 - Z=20)

First Ionization Energy vs. Atomic #

Group Trends in Ionization Energy ● In general, first ionization energy decreases from top to bottom within a group o Recall, atomic size increases as atomic number increases in a group

Group Trends in Ionization Energy Atomic size increases → Nuclear charge has smaller effect on e- in highest occupied energy level → Less energy require to remove an e- from this energy level → Lower first ionization energy

Periodic Trends in Ionization Energy ● In general, first ionization energy of representative elements tends to increase from left to right across a period

Periodic Trends in Ionization Energy Nuclear charge increases across the period, but shielding effect remains constant → Increase in nucleus’ attraction for an e→ More energy to remove an e- from the atom → Higher first ionization energy

Example Which of the elements in each pair has a larger ionization energy? Explain. 1) sodium or potassium? 2) magnesium or phosphorus?

Example Would you expect positive ions to be bigger or smaller than parent atom? Explain. Would you expect negative ions to be bigger or smaller than parent atom? Explain.

Trends in Ionic Size ● Cations always smaller than parent atom o Lose e-s  Attraction between remaining e-s and nucleus increases ● Group A metals tend to lose all their outermost e-s during ionization o Ion has one fewer occupied energy level

Trends in Ionic Size ● Anions always larger than parent atom o Gain e-s  Attraction of the nucleus for any one e- decreases

Trends in Ionic Size ● From left to right across a period: o Gradual decrease in the size of cations o Gradual decrease in the size of anions ● Increases within groups

Trends in Electron Affinity Electron affinity (EA): ● The energy changes that accompanies the addition of an electron to a gaseous atom

Trends in Electron Affinity ● Most elements release energy when they gain an electron o Most electron affinities are negative ● Show relative ease by which atoms gain electrons o Positive EA -- energy is absorbed o Negative EA -- energy is released

Trends in Electron Affinity

Trends in Electron Affinity ● Not as easy to determine experimentally ● Trends in electron affinity are unclear o Generally increases left to right across period  Atoms are smaller, nuclear charge increases o Generally decreases top to bottom down group  Atoms are bigger

Example In general, would you expect nonmetals to have larger electron affinities than metals? Explain.

Trends in Electronegativity Electronegativity: ● The ability of an atom of an element to attract e-s when an atom is in a compound ● Scientists use ionization energy to calculate electronegativity

Trends in Electronegativity ● Expressed in units called Paulings (named after Linus Pauling who defined electronegativity)

Trends in Electronegativity ● In general, electronegativity: o Decreases from top to bottom within a group o Increases from left to right across a period (for representative elements)  Metals (far left) have low value  Nonmetals (far right) have high values  Irregular among transition metals  None for noble gases

Trends in Electronegativity ● Least electronegative: cesium o Tends to lose e-s to form positive ion ● Most electronegative: fluorine o Tends to attract shared e-s or form negative ion

Summary of Trends ● The trends that exist among these properties can be explained by variations in atomic structure o Increase in nuclear charge within groups and across periods explains many trends o Within groups, an increase in shielding has a significant effect