Chemistry A Molecular Approach AP Edition Chapter 8

Chapter 8: Periodic Properties of the Elements 

• In 1869, Dmitri Mendeleev found a pattern in the elements. He organized them according to these patterns  
• Quantum mechanics explains how electrons arrange in an atom 

8.1 Nerve Signal Transmission

• Sodium and Potassium are being pumped through your body. But the ions travel in opposite directions. The sodium is pumped out of cells while potassium is pumped in 
• When ion channels open, the ions flow downwards 
• The movement of ions is what causes nerve signals to be sent to the brain 
• Periodic properties are predictable based on the element's position on the table 

8.2 The Development of the Periodic Table 

• The modern periodic table was created by Russian chemist Dmitri Mendeleev 
• Elements are ordered by increasing mass and certain properties repeat periodically. Similar properties fall in the same column 
• Mendeleev could predict elements that hadn't been discovered yet with this table 
• Listing elements by atomic number rather than mass was more successful 

8.3 Electron Configurations: How Electrons Occupy Orbitals

• The quantum-mechanics theory describes the behavior of elements 
• Electron configurations shows the particular orbital that electrons occupy 
• The ground state is the lowest energy state 
• How electrons occupy orbitals also depends on the spin of the electron and the sublevel energy splitting 
• The spin determines how many electrons can be in one orbital 
• Sublevel energy splitting determines the order of orbital filling 

Electron Spin and the Pauli Exclusion Principle 

• An orbital diagram symbolizes the electrons as an arrow and the orbital as a box 
• The direction of the arrow represents the electrons spin 
• The spin can either be up (Ms = +1/2) or down (Ms = -1/2) 
• The Pauli exclusion principle states that no two electrons can have the same four quantum numbers. This means the electrons cannot have the same spin in the same orbital box. The other three quantum numbers must be the same because the electrons are in the same orbital  

Sublevel Energy Splitting in the Multielectron Atoms 

• E(s orbital) < E(p orbital) < E(d orbital) < E(f orbital) 
• Coulomb's Law is the attractions and repulsions between charged particles 
• Attraction decreases as particles get farther apart
• Charges can repel each other 
• Opposite charges attract each other 
• Magnitude increases as charge increases 
• Shielding protects certain electrons from the full effects of nuclear charge 
• The electrons that is being shielded is experiencing effective nuclear charge 
• Penetrations is when an electron gets closer to the nucleus 
• Penetration causes an electron to feel a greater nuclear charge and have a lower energy 
• The splitting of energy sublevels deals with the distribution of electrons 
• 2p orbitals are farther from the nucleus than 2s orbitals 

Electron Configurations for Multielectron Atoms 

• The aufbau principle says that only two electrons with opposing spins are allowed in each orbital 
• Hund's rule says that you must fill each orbital before putting two electrons in the same orbital 
• When two electrons are in the same orbital, they repel each other slightly 
• Lower energy levels fill before higher energy levels 
• Noble gases can be used to represent inner electrons 

8.4 Electron Configurations, Valence Electrons, and the Periodic Table 

• Elements with similar properties are in the same column 
• Valence electrons are important for chemical bonding 
• Elements in the same column have the same number of valence electrons 
• Valence electrons are on the outer shell 
• All other electrons are referred to as core electrons 

Orbital Blocks in the Periodic Table

• The first two columns of the periodic table fill the s orbitals
• The next six columns fill the p orbitals 
• Transition elements make up the d orbitals 
• Lanthanides and actinides form the f orbitals
• The number of columns represents the number of electrons that can occupy the sublevel of that block 
• Helium has 2 valence electrons 
• Otherwise, the amount of valence electrons is equal to the lettered group number 

Writing an Electron Configuration for an Element from Its Position in the Periodic Table 

• The organization of the periodic table makes us able to write the electron configuration for any element 
• Inner electron configuration is equal to the previous noble gas 
• Outer electron configuration can be counted 

The Transition and Inner Transition Elements

• Lanthanides and actinides are sometimes called inner transition elements
• 4s orbitals are lower in energy that 3d orbitals 
• Orbitals fill from left to right   

8.5 The Explanatory Power of the Quantum-Mechanical Model 

• The chemical properties of an element are generally determined by their number of valence electrons 
• Elements within a column have the same number of valence electrons
• Elements within a column have similar properties 
• Atoms with 8 valence electrons, and 2 in the case of Helium, are stable 
• Noble gases are the most stable and unreactive family of the Periodic Table 
• Alkali metals are very reactive because they can easily lose one electron to have a noble gas configuration 
• Alkaline earth metals are also reactive 
• Forming cations always requires energy 
• Halogens are the most reactive nonmetals 
• Halogens are reactive because they can easily gain one electron to have a noble gas configuration 

8.6 Periodic Trends in the Size of Atoms and Effective Nuclear Charge 

• The volume of an atom is largely made up of electrons 
• One way to determine the radii of an atom is to find the distance between nonbonding atoms 
• Nonbonding atomic radius is also called the van der Waals radius 
• The bonding atomic radius or covalent radius is also used to determine the size of an atom 
• For nonmetals the bonding radius is one-half the distance between two of the atoms bonded together 
• For metals the bonding radius is one-half the distance between two of the atoms next to each other in a crystal 
• The atomic radius os always smaller than the nonbonding atomic radius 
• The atomic radius increases moving down a column 
• The atomic radius decreases moving across a period (row) 
• The atomic radius is largely determined by the number of valence electrons  

Effective Nuclear Charge 

• Moving across a period, the atomic radius is determined by the inward pull of electrons from the nucleus 
• The average charge experienced by an electron is the effective nuclear charge 
• The effective nuclear charge can be found by subtractive the charge of the shielded electrons from the actual charge 
• Core electrons shield outer electrons from nuclear charge 

Atomic Radii and the Transition Elements 

• Most transition elements do not follow the same trend as main group elements 
• When moving across a row, the radii of transition elements stays constant 
• This is because the number of outermost energy level is constant and the number of outermost electrons stays constant 
• Because the number of electrons stays constant, the elements have the same effective nuclear charge, making the radius constant 

8.7 Ions: Electron Configurations, Magnetic Properties, Ionic Radii, and Ionization Energy 

Electron Configurations and Magnetic Properties of Ions

• For anions, add the number of electrons indicated by the charge
• For cations, subtract the number of electrons indicated by the charge 
• For transition metals, remove electrons in the highest n0value orbital
• V: [Ar] 4s^2 3d^3
• V^2+ : [Ar] 4s^0 3d^3
• When a atom contains unpaired electrons, it is paramagnetic 
• At atom that has every electron paired is diamagnetic 

Ionic Radii

• Cations are much smaller than their normal atoms 
• This is because one of the outermost electrons is lost, allow the other electrons to be closer to the nucleus
• Anions are much larger than their normal atoms 
• This is because more electrons cause more repulsion, making the atom grow 

Ionization Energy 

• The ionization energy of an atom is the energy required to remove an electron from the atom (in the gas state) 
• Ionization energy is always positive 
• The energy required to remove a second electron is called the second ionization energy 
• The energy required to remove a third electron is called the third ionization energy 
• The second ionization energy is NOT the energy required to remove two electrons

Trends in First Ionization Energy

• Electrons in the outermost shell are very far away from the nucleus, therefore they are held less tightly 
• Ionization energy lowers going down a column 
• Ionization energy increases going right across a row
• Ionization energy increases across a row because there is a greater effective nuclear charge  

Exceptions to Trends in First Ionization Energy 

• Exceptions are caused by electrons is the s and p blocks 
• 2p orbitals have higher energies and their electrons are easier to remove 
• There is also an exception between Nitrogen and Oxygen. This is because there is a repulsion in the 2p block of Oxygen, making that electron easier to remove

Trends in Second and Successive Ionization Energies  

• The first ionization energy involves removing valence electrons, this makes the IE a lower value 
• The second ionization energy requires the removal of a core electron from a noble gas configuration, making the IE2 very high 
• The third ionization energy require removing an electron that is experiencing effective nuclear charge which also requires a lot of energy 

8.8 Electron Affinities and Metallic Character

• Electron affinity is how easily an atom will accept an addition electron 
• Electron affinity is important to bonding because it involves the transfer/sharing of electrons 
• Metallic character is important because there are many metals on the periodic table 

Electron Affinity 

• Electron affinity of an atom is the energy change associated with the gaining of an electron 
• Electron affinity is usually negative 
• The trend of electron affinity going down a column is not regular 
• In group 1A elements, electron affinity becomes more positive moving down the column 
• For main group elements, electron affinity generally becomes more negative going right across a row because an electron is added 

Metallic Character

• Metals are good conductors of heat and electricity 
• Metals are malleable and ductile 
• Metals are shiny and lose electrons in chemical reactions 
• Nonmetals have varied physical properties but tend to be poor conductors of heat and electricity and gain electrons in chemical reactions 
• Elements on the left side of the periodic table are more likely to lose electrons
• Elements on the right side of the periodic table are more likely to gain electrons 
• Moving right across a row, metallic characteristics decrease 
• Moving down a column, metallic characteristics increase 

8.9 Some Examples of Periodic Chemical Behavior: The Alkali Metals, the Halogens, and the Noble Gases 

• Alkali metals can easily lose an electron to have noble gas configurations 
• Alkali metals are very reactive 
• The halogens can easily gain an electron to have noble gas configurations 
• The halogens are the most reactive nonmetals
• Noble gases have full electron configurations 
• Hydrogen acts like a nonmetal and is not included in 1A metals 

The Alkali Metals (Group 1A) 

• Atomic radius increase and ionization energy decreases going down the 1A column 
• Density increases (except K) going down 
• The mass increases 
• Melting points are low for metals and the melting point decreases going down 
• Alkali metals are good reducing agents 
• Alkali metals exist naturally in their oxidized state 
• Reactivity of alkali metals increases going down 

The Halogens (Group 7A) 

• All halogens are powerful oxidizing agents 
• Fluorine is the most powerful oxidizing agent 
• Halogens react with metals to form teal halides 
• Halogens tend to gain electrons 

The Noble Gases (Group 8A) 

• All noble gases are gases at room temperature 
• Noble gases are unreactive
• They are inert