Periodic Table SPW 234 Chapter 19

Let’s Review! Chemical properties Any property that can only be tested by changing the chemical make-up of the substance. Physical properties Any property that can be tested without changing the chemical make-up of the substance Atomic mass Mass of protons and neutrons Atomic number Unique to each element, same as number of protons

Dmitri Mendeleev 1870: 63 elements known to man He organized them in order of their atomic mass, and saw a pattern from their properties. Was working on this while Thomson and Rutherford were still “exploring” the atom

Dmitri Mendeleev Arranged his table with repeating properties in columns, starting a new row each time the chemical properties repeated Left blank spaces in his table, concluding that these spaces were elements that hadn’t been discovered yet. Based on the patterns and the other elements around the blank space, he predicted the properties of those elements

What he called ekasilicon – it was discovered a few years later An Example What he called ekasilicon – it was discovered a few years later Prediction Atomic Mass 72 amu Density 5.5 g/mL Appearance dark gray metal Melting Point high melting point Germanium 72.6 amu 5.3 g/mL gray metal 937o C

Mendeleev’s Table

Henry Moseley Mendeleev’s table worked because as protons increase, atomic mass should increase, but if there are fewer neutrons it could decrease Errors arose because he was arranging the table with the wrong number 1910: Discovered atomic number and rearranged the periodic table using this number, it fell into perfect order

Periodic Law Periodic Law: physical and chemical properties of the elements are periodic functions of their atomic numbers In other words, when the elements are arranged by their atomic numbers, you should see chemical and physical properties repeating themselves

Rows Left to right – called periods Elements in the same periods show patterns left to right - conductivity/reactivity change, elements become less metallic The period # indicates the number of energy levels each atom in the row has

Columns Top to bottom – called groups Elements in a group have similar chemical properties and show trends top to bottom The elements in the same group (column) have the same number of valence electrons

Valence Electrons The trends found in the periodic table are a result of electron arrangement, specifically, the number of valence electrons Valence Electron: electrons in the outermost energy level

Valence Electrons The group number of an element will tell you the number of valence electrons it has Group 1: 1 valence electron Group 2: 2 valence e- ’s Skip 3-12 Group 13: 3 valence e- ’s Groups 14-18: 4, 5, 6, 7, and 8 valence e- ’s respectively.

Ion When a neutral atom gains/loses electron(s) through bonding, the atom is no longer neutral and has an excess charge It becomes an ion Ion: a charged atom

Ion All atoms want 8 valence electrons in the outer shell – this makes the shell full and the atom stable Elements close to having 8 tend to be the most reactive. Elements already “full” are considered inert, they don’t react because they don’t need to The number of electrons an atom can gain or lose is equal to the number of valence electrons it has

Let’s Practice Group 16 Give Up? or Gain? Group 13 Give Up? or Gain?

Ion Protons = positive charge Electrons = negative charge p+ # CANNOT change, but e- # can So… If an atom GAINS electrons, is it more positive, or negative? If they LOST electrons?

ION Cation: atoms that LOSE electrons, becoming more positive Anion: atoms that GAIN electrons, becoming more negative

ion How do we know if an atom is an ion? Cations have a +, and anions have a – superscript If an atom has gained 3 electrons It has 3 MORE negative particles than positive particles, it is more negative = Al3- If an atom has lost 3 electrons It has 3 LESS negative particles than positive particles, it is more positive = Al3+

The Periodic Table Divided into three major categories based on general properties Metals, Nonmetals, Metalloids (semiconductors)

Metals Like to give up valence electrons Physical Properties: high luster (shiny), conductive (heat and electricity), malleable (bendable), ductile (stretchable), high density, high melting point Chemical Properties: Most will react with oxygen

Nonmetals Like to gain electrons Physical Properties: dull, don’t conduct, brittle, low density, low melting points Can be solid, liquid or gas at room temperature depending on the element.

Metalloids (Semiconductors) Share properties of both metals and nonmetals Can be shiny or dull, conduct ok, ductile and malleable or brittle These elements have become really important because of the computer revolution Computer chips are made out of semiconductors (normally Si)

Families The periodic table can be further broken down into families Families of elements have similar properties because they have the same number of valence electrons

families Metals: Groups 1-12 Alkali Metals, Alkaline-Earth Metals, Transition Metals Groups 13-16 contain both nonmetals/metalloids Nonmetals: Oxygen, Nitrogen, Carbon, Sulfur, Phosphorus, and Selenium Metalloids: Boron, Silicon, Germanium, Arsenic Antimony, and Terellium The group is named by the first element in the column Nonmetals: Groups 17-18 Halogens, Noble Gases

Hydrogen Hydrogen is in group 1 but is not an alkali metal, because it is only 1 proton and 1 electron (no neutrons) Its properties are closer to a nonmetals than to a metal it is a colorless, odorless, explosive gas with oxygen

Group 1: Alkali Metals (excluding H), 1 valence e- Very reactive, especially with water Soft, shiny white metals (can be cut with a knife!)

Group 2: Alkaline-Earth Metals 2 valence e- Not as reactive as alkali, but still very reactive. Magnesium is used in flash bulbs

Groups 3-12: Transition Metals 1, 2, 3, or 4 valence e- Most are silver and not that reactive so they have more everyday uses.

Groups 3-12: Transition Metals Two bottom rows, or innertransition metals Lanthanide Series: also called rare- earth metals Actinide Series: very radioactive and not easily found in nature

Groups 13-16 Boron Group: Group 13, 3 valence e- Aluminum is most common and abundant element on the planet. Carbon Group: Group 14, 4 valence e- Pure carbon can be diamonds, soot, or graphite, silicon and germanium are used for computer chips

Groups 13-16 Nitrogen Group: Group 15, 5 valence e- Nitrogen makes up 78% of the air, Phosphorus is in soaps, and Arsenic is a well known poison Oxygen Group: Group 16, 6 valence e- ’s Oxygen makes up 21% of the air and is necessary for things to burn

Group 17: Halogens 7 valence e- All nonmetals (can be solid, liquid or gas) Extremely reactive with alkali metals “Chlorine” added to pools as a disinfectant is a compound containing Chlorine, by itself chlorine is a green gas

Group 18: Noble Gases 8 valence e- ’s (except Helium) Full outer shell of electrons All are gases and extremely non- reactive (inert) and found in the atmosphere “Neon” lights contain a variety of Noble Gases