Physical Setting/Chemistry

Physical Setting/Chemistry Classifying the Elements and Trends in the Period Table Br. Jabreal

Aim: What categories of Elements are Present in the Periodic Table? • Do Now: List the elements in group 18. Are these metals, nonmetals or metalloids?

Noble Gases • Group 18: He, Ne, Ar, Kr, Xe, Rn are known as the Noble Gases • They are stable and unreactive – they almost always will never react with other elements • For that reason, they are often also called the inert gases • Their outermost energy level is always full with electrons

Halogens • Group 17: F, Cl, Br, I, At are known as the Halogens • Are these metals or nonmetals? • These are nonmetals • React with group 1 elements to form a class of compounds called “salts” • E.g. NaCl • “Halogens” comes from latin“hals” meaning salt and “genesis” meaning to be born.

Alkali Metals • Group 1 elements are called “Alkali Metals.” • Includes … • Li, Na, K, Rb, Cs, Fr, but not H

Alkaline Earth Metals • Group 2 metals are known as the alkaline earth metals

Transition Metals • Middle part of the Periodic Table • Group 3 to Group 12 • Give 3 examples of transition metals

Electron Configurations in Groups • Elements divided into their groups based on electron configurations • How many electrons are in the outermost shell of halogens? Of noble gases? Of Alkali metals? Of Alkaline earth metals? • Rule for non-transition metals: Group number = number of electrons in outermost energy shell (count groups 13-18 as 3A - 8A).

Aim: What are the patterns in the periodic table? • Do Now: What is indicated by elements in the same group? What is the significance of periods?

Trends in Atomic Size • Atoms don’t have sharply defined boundaries, making measuring an atoms size difficult • One way is by measuring the distance between the nuclei of identical atoms connected to each other

Atomic radius is one half the distance between the nuclei of two atoms of the same element Ex. Atomic radius of Hydrogen using two H atoms

Trends in Atomic Size • What happens to atomic size as you go across a period (left to right)? • Let’s figure it out. Pick a period. • Check table S for the atomic radius of the first element in the period • Check Table S for the atomic radius of the next few elements • What’s the pattern? • Atomic size decreases from left to right across a period

Trends in Atomic Size • What happens to atomic size as you go down a group? • Let’s figure it out. Pick a group. • Check Table S for atomic radius of element at top of the group. • Check Table S for atomic radius of next few elements. • What’s the pattern? • Atomic size increases from top to bottom within a group

Group Trends in Atomic Size • Why does atomic size increase down the group? • Two things happen going down the group: • 1. Nuclear positive charge from more protons increases • 2. More electrons and more occupied energy level shells • But the extra energy shells shield electrons from pull of the protons, making the radius bigger

Periodic Trends in Atomic Size • Why does atomic size decrease left to right across a period? • How does each element compare to the preceding element in terms of protons and electrons? • Each has one more proton and one more electron • Same occupied energy levels, so no shielding effect … but more pull from the increased protons and the nuclear charge • Causes size of atom to decrease

Ions • Ion: An atom or group of atoms that has a positive or negative charge • We said that generally, number of protons = number of electrons • However, positive and negative ions form when electrons are transferred between atoms

Metallic elements, such as sodium, tend to form ions by losing one or more electrons. • What happens to the charge on these atoms? • They become cations: ions with a positive charge • Because there are less negatively charged electrons than positively charged protons

Atoms of nonmetallic elements, such as chlorine, tend to for ions by gaining one or more electrons • How would this affect the atom’s charge? • The atoms become anions: Ions with negative charges. • Because there are more negatively charged electrons than positively charged protons

Do Now: What is an ion? Summarize the types of ions and how they form.

Ionization Energy • Recall that electrons can move energy levels when they absorb energy • Sometimes there is enough energy to overcome attraction of the protons in the nucleus • Ionization Energy: Energy required to remove an electron from an atom • First ionization energy is for removal of 1st electron • Second ionization energy is for removal of an additional (second) electron

Trends in Ionization Energy • Let’s discover the trends within groups and periods using the Periodic Table and Table S again … • First ionization energies tend to decrease from top to bottom within a group and increase from right to left across a period

Group Trends in Ionization Energy • Why does ionization energy tend to decrease from top to bottom within a group? • As you go down the group, atomic size increases • As the atomic size increases, the nuclear charge has a smaller effect on the highest occupied energy level. • Therefore, less energy required to remove one electron (lower ionization energy)

Periodic Trends in Ionization Energy • Why does ionization energy increase from left to right within a period? • Nuclear charge increases from more protons as you go across the period, from left to right… And shielding effect of energy levels remains constant as highest energy shell of electrons stays the same • Increasing strength of pull from nucleus • Therefore, it requires more energy to remove an electron (higher ionization energy).

Trends in Ionic Size • Metals tend to lose electrons and nonmetals tend to gain electrons • Cations (positive ions) are always SMALLER than the atoms from which they form • Anions (negative ions) are always LARGER than the atoms from which they form

Electronegativity • Electronegativity is a property that determines the type of bonds that atoms will form. • Electronegativity: The ability of an atom of an element to attract electrons when the atom is in a compound • Measured in units called Paulings • Fluorine is the most electronegative element • Values for Noble Gases omitted because they do not form many compounds

Trends in Electronegativity • In general: • Electronegativity decreases going down a group • Electronegativity increases going across a period

Summary of Trends

Summary of Trends • Why do the trends exist? • Trends that exist among these properties can be explained by variations in atomic structure • Increase in nuclear charge within groups and across periods • Within groups an increase in shielding has a significant effect

Practice • As you move up and to the right on the periodic table: • Atomic radius increases and electronegativity increases • Atomic radius decreases and electronegativity increases • Atomic radius increases and electronegativity decreases • Atomic radius decreases and electronegativity decreases

Practice • Atoms and ions with the same number of electrons are called isoelectronic. • Write the symbol for a positive ion and a negative ion that are isoelectronic with krypton.

Homework • Read Section 6.3 • Section 6.3 Assessment • Questions 21, 22, 23

Physical Setting/Chemistry