Section 1 Electron structure and the periodic table Periodic Table II Section 1 Electron structure and the periodic table

Recall that the Periodic table is arranged according to properties? From reactive metals sodium Mercury To less reactive metals

Recall that the Periodic table is arranged according to properties? helium …and inert (noble) gases non-metals silicon To semi-metals carbon

The Periodic table is also arranged according to electron arrangement 8 Valence electrons 1, 2, 3 Valence electrons 4 to 7 Valence electrons

So the properties must depend on electron arrangement! The Periodic table is also arranged according to electron arrangement So the properties must depend on electron arrangement!

Periods (horizontal) 1, 2, 3 etc. Periods have elements with same number of occupied energy levels (shells) OR Have the same outer (valence) shell

Periods (horizontal) 1, 2, 3 etc. Periods have elements with same number of occupied energy levels (shells) OR Have the same outer (valence) shell Example: Elements in period 3 have electrons in three energy shells OR the 3rd shell is the valence level. Al e- config. = 2 – 8 – 3

Periods (horizontal) 1, 2, 3 etc. Periods have elements with same number of occupied energy levels (shells) OR Have the same outer (valence) shell Within a period electrons vary from 1 to 8 Valence electrons Also, the properties change as the atomic number increases

Groups (vertical) 1,2,…18 8 7 Valence electrons 6 1 2 5 4 3 1 or 2 Groups have elements with the same number of valence electrons 8 7 Valence electrons 6 1 2 5 3 4 1 or 2 Elements with the same number of valence electrons make for similar properties in groups

Crash course chemistry Click here for Crash course chemistry on Periodic Table

Test Your Understanding What aspect of an atoms structure allows you to predict its properties? How does the arrangement of elements on the table support your conclusion? Why do the transition elements have such similar but unique properties? What change in the atomic structure of the elements in a period accompanies the change in properties we observe?

HDYK?

Periodic Trends > Size of atoms > Strength of atoms Section 2 Periodic Trends > Size of atoms > Strength of atoms

2A) Periodic Trends: Size of atoms Atoms get larger down the table How can you explain these trends? (how do the structure of atoms differ? But get smaller left to right

Firstly, how are atoms measured? Shooting X-rays at a sample reveals where the nuclei are located, but not the electrons. Atomic nuclei

Firstly, how are atoms measured? If you assume that the outer shell of each atom reaches half way between two atoms. Then ½ the distance between the two nuclei is each atom’s Radius. D/2 radius radius

Secondly, what determines the radius? The size of an atom depends on two factors: #1 The number of electron shells. More shells take up more space Atoms get bigger!

Periodic Trends: Size of atoms As atomic number increases going down a group each atom has an additional energy level, further from the nucleus. Radius goes UP since the valence shell is now further from the nucleus.

The pull from the protons in the nucleus Secondly, The size of an atom depends on two factors: #2 The pull from the protons in the nucleus

Secondly, what determines the radius? The size of an atom depends on two factors: A stronger nuclear charge will pull electrons closer. so atoms are smaller.

As Atomic number increases across a period, atoms are gaining more protons in their nuclei (greater nuclear charge) Radius decreases - since the extra positive charge from the protons pulls the electrons closer to the nucleus 11+ 12+ 13+ 14+ 15+ 16+ 17+ Large Small

Summary: Atomic radii Click here for smaller Khan academy tutorial on atomic radius smaller Overall Trend larger

Test your learning How is the size of an atom expressed? Why? Which atom is larger; H or He? Why? Which is larger; and argon atom or a potassium atom? Why? Describe the overall trend in atomic radius for elements on the periodic table.

Introduction to chemical reactions Atomic “strength” Section 2B Introduction to chemical reactions Atomic “strength”

Why do atoms react with one another? Recall that the group 18 Noble gases have 8 valence electrons? Recall that the group 18 Noble gases are very stable? Atoms with 8 valence electrons are very stable since they require a lot of energy to remove any of their electrons.

Why do atoms react with one another? Nonmetal atoms like oxygen can become stable by gaining some additional electrons. The new ion that is formed is no longer neutral

Why do atoms react with one another? Ions like this now called “oxide “ are negative since they have two more electrons than protons Negative ions like this are called an-ions

Why do atoms react with one another? Its new symbol is O2- Its new electron configuration is now 2-8 like neon Negative ions like this are called an-ions

Why do atoms react with one another? Metallic atoms like sodium can become stable by losing some electrons.

Why do atoms react with one another? Sodium’s electron configuration was 2-8-1 Now becomes 2-8 Stable like neon.

Why do atoms react with one another? The symbol for the sodium ion is Na1+ since it lose an electron it now has less electrons than protons and carries an excess positive charge. Positive ions are called Cat-ions

Forming an ion: So atoms lose or gain electrons to form ions Why? O 2- = 8 P+ + 10 e- = -2 charge Na+ = 11 P+ + 10 e- = + 1 charge But, why are some atoms losers, and other atoms gainers?

Atomic Strength: Ionization Energy Energy is required to remove an electron from an atom. This is called it ionization energy (energy to ion-ize) Mg + 738 kJ ---> Mg+ + e- Energy Notice energy is added to Mg to remove one electron. This is called the FIRST ionization energy because we removed only one of the outer electrons ion and electron

Atomic Strength: Ionization Energy Energy is required to remove an electron from an atom. This is called it ionization energy (energy to ion-ize) Mg+ + 1451 kJ ---> Mg2+ + e- It takes even more energy to remove a second electron. This is the SECOND ionization energy.

Atomic Strength: Ionization Energy Energy is required to remove an electron from an atom. This is called it ionization energy (energy to ion-ize) Metals like magnesium will lose electrons since their ionization energies are relatively low.

Atomic Strength: Ionization Energy Energy is required to remove an electron from an atom. This is called it ionization energy (energy to ion-ize) In other words, metals are “weak”

Atomic Strength: Ionization Energy Energy is required to remove an electron from an atom. This is called it ionization energy (energy to ion-ize) F + 1681kJ ---> F+ + e- Removing an electron from a nonmetal requires too much energy

Atomic Strength: Ionization Energy Energy is required to remove an electron from an atom. This is called it ionization energy (energy to ion-ize) F + 1681kJ ---> F+ + e- Nonmetals like fluorine are too strong to lose electrons

Atomic Strength: Ionization Energy Energy is required to remove an electron from an atom. This is called it ionization energy (energy to ion-ize) F + 1681kJ ---> F+ + e- In fact, nonmetals gain electrons from metals and become stable.

Atomic Strength: Ionization Energy Energy is required to remove an electron from an atom. This is called it ionization energy (energy to ion-ize) F + e-  F- + 328 kJ In fact energy is released when F gains and electron

Atomic Strength: Ionization Energy Energy is required to remove an electron from an atom. This is called it ionization energy (energy to ion-ize) F + e-  F- + 328 kJ This is called electron affinity

Trends in Ionization Energy Ionization energy increases across a period because the positive nuclear charge increases (more protons). losers 11+ 12+ 13+ 14+ 15+ 16+ 17+ Period 3 496 kJ 1012 kJ 1251 kJ As the nuclear charge gets stronger, they exert a greater pull on the outer electrons

Trends in Ionization Energy Ionization energy increases across a period because the positive nuclear charge increases (more protons). Ionization energy increases gainers losers Metals (left side) are weak - lose electrons more easily: losers Nonmetals (right side ) strong - don’t lose electrons: gainers

Going down a group Atoms get larger: this means there is a greater distance from the nucleus to valence electrons. So they get weaker. and Ionization energy decreases. decreases Ionization energy

Atoms with lots of inner shells. Ionization energy decreases Shielding Effect Atoms with lots of inner shells. Ionization energy decreases *Each additional energy level adds electrons that repel the valence electrons

Notice the Overall Trend Like radius, ionization energy follows the diagonal trend.

Non-metal atoms are strong And small. Notice the Overall Trend Metal atoms are weak And large

Periodicity: a trend that occurs at regular intervals Click here for Khan academy Tutorial Ionization energy trends Notice how ionization energy increases, then decreases with each new period.

What does diagram this illustrate?

Learning check What is ionization energy? What does it indicate about an element? Which will have a higher ionization energy H or He? Why? Which will have a larger ionization energy Ar or K? Why? Describe the overall trend in ionization energy on the periodic table.

Other important trends Section 3 Other important trends

Ion Sizes Metal ions are SMALLER than the atoms from which they come. Forming a cation. Metal ions are SMALLER than the atoms from which they come. Loss of entire energy level so size DECREASES.

Forming an anion Nonmetal ions are LARGER than the atoms from which they come. Additional electrons repel each other more so size INCREASES (“swells up”)

Think about it Which of the following is the smallest: O2- ion, Ne atom, or Mg2+ ion? Why? Hint: Each has 10 electrons. But: How many protons? +10 +8 +12 oxide Neon magnesium 8 protons 10 protons 12 protons

Think about it Which of the following is the smallest: O2- ion, Ne atom, or Mg2+ ion? +8 +10 +12 Largest    middle    smallest oxide Neon magnesium 8 protons 10 protons 12 protons As nuclear charge increases, the size will decrease

Test your understanding What is the electron configuration for a magnesium atom? In terms of electrons, what will likely happen to magnesium during a chemical reaction? What will be its new electron configuration? Draw the Bohr diagram for the magnesium atom and the magnesium cation. Which will be larger? Why?

Atomic # 12 = Mg 9 An atom of an element has a total of 12 electrons. An ion of the same element has a total of 10 electrons. Which statement describes the charge and radius of the ion? (1) The ion is positively charged and its radius is smaller than the radius of the atom. (2) The ion is positively charged and its radius is larger than the radius of the atom. (3) The ion is negatively charged and its radius is smaller than the radius of the atom. (4) The ion is negatively charged and its radius is larger than the radius of the atom. 12 P+ + 10 e-

Loser metal

Chemical strength: Electronegativity measure of the ability of an atom in a molecule to attract electrons to itself. Attraction for a pair of electrons shared in a chemical bond  Chemical bond 

Electronegativity: same as trends in Ionization energy Why no values for group 18? Increases: stronger nucleus = stronger force Decreases: More Energy Levels = Weaker force decreases increases

Why? HDYK? Which has a higher: 1st ionization energy? Mg or Ca ? Al or S ? Cs or Ba ? Circle one in each pair Why? Which is more electronegative? F or Cl ? Na or K ? Sn or I ? HDYK?

Ionization Energy

Introduction to Chemical Reactivity Atoms lose or gain electrons to become stable like group 18 elements (8 valence e- ‘s) Ex: Ne 2 - 8 What’s the magic number?

What elements sit beyond group 18? Metals lose To become Like Noble gases Nonmetals gain To become like Noble gases F: 2-7, becomes 2-8 Na: 2-8-1, becomes 2-8

Metals – group 1,2 and transitional, etc. Chemically weak, tend to lose outer electrons to stronger atoms Ex: Mg 2 - 8 - 2 loses 2e- becomes 2 - 8 like Ne Becomes a positive ION Mg2+ ION = charged “atom” CATION – a positively charged ion

gain 1 e- Nonmetals – groups 15, 16, 17 Chemically strong tend to gain electrons from weaker atoms Ex: Fluorine atom F 2 - 7 gains 1 e- becomes 2 - 8 like Ne Becomes a negative fluoride ION: F1- ANION - a negative ion

13 When a lithium atom forms an Li+ ion, the lithium atom (1) gains a proton (2) gains an electron (3) loses a proton (4) loses an electron 37 What is the total number of electrons in a Cu+ ion? (1) 28 (3) 30 (2) 29 (4) 36

F gains e-

Activity / Reactivity More Active metals F More Active nonmetals Fr Metals –”losers” must lose electrons during reactions weakest (lowest electronegativity) are most active (francium) Nonmetals – gainers must gain electrons during reactions strongest (highest electronegativity) are most active (fluorine) More Active metals F More Active nonmetals Fr

Metallic / nonmetallic “character” More (loser) metallic character More (gainer) Nonmetallic character

33 As the elements in Group 17 on the Periodic Table are considered from top to bottom, what happens to the atomic radius and the metallic character of each successive element? (1) The atomic radius and the metallic character both increase. (2) The atomic radius increases and the metallic character decreases. (3) The atomic radius decreases and the metallic character increases. (4) The atomic radius and the metallic character both decrease.

Hint: standard temperature = 00C For questions 73 – 76 Refer to the table below and your knowledge of chemistry Hint: standard temperature = 00C

181 98 64

1) How many elements are solids at STP? 2) Which element is a noble gas? 3) Letter Z below corresponds to a different element. Elements G,Q, L and Z are in the same group on the periodic table as shown to the right. 4) Based on the trend in melting points for G, Q and L, estimate the melting point of Element Z

Can you explain this trend in ionization energy? +19 1) K + 419 kJ ---> K+ + e- Energy 2) Na + 496 kJ ---> Na+ + e- Energy +11 3) Mg + 738 kJ ---> Mg+ + e- Energy 4) Mg+1 + 1451 kJ ---> Mg+2 + e- Energy +13 Hint: compare the structures on the right 