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PERIODICITY Chapter 6 Mrs. Medina 1 Searching for an organizing principle • Year 1700, only 13 elements were known – Copper, silver, gold, iron… • Chemists began using scientific methods and the rate of discovery increased. • How would chemists know if all the elements had been discovered? – They had to be organized into some sort of system. • Chemists used the properties of elements to sort them into groups. 2 Mendeleev’s Periodic Table • In 1869, a Russian chemist and teacher, Dmitri Mendeleev published a periodic table. • He arranged the elements in order of increasing atomic mass. – By doing this, the elements fell into groups because the properties began to repeat themselves. • This repetition of properties allowed Mendeleev to predict the elements gallium and germanium earning him international fame for his table. 3 The Periodic Law • Mendeleev’s table placed the elements in order of increasing atomic mass but this arrangement had to be broken sometimes because otherwise the properties would not match up. • Moseley discovered the proton in 1913. • When the elements were placed in increasing atomic number, all the repeating properties fell into place without any problems. • In the modern periodic table, elements are arranged in order of increasing atomic number. 4 The Periodic Law • Periodic Law: When elements are arranged in order of increasing atomic number, there is a repetition of their physical and chemical properties. – Elements within a group (column) have similar properties – Properties in a period (row) change as you move from left to right. 5 Metals, NonMetals and Metalloids 6 Metals • Most elements are metals • Good conductors of heat and electricity • High luster, or sheen b/c metals reflect light • All metals are solids at room temperature except one: mercury (Hg) which is a liquid. • Ductile – draw into wires • Malleable – hammer into thin sheets 7 Metals 8 Nonmetals • Include the noble gases and hydrogen • Among nonmetals there is a greater variation in physical properties – Most are gases, a few are solids, and one is a liquid. • They have properties that are opposite of the metals: – Poor conductors of heat and electricity • carbon being an exception – Solid nonmetals will be brittle (shatter easily) 9 Nonmetals 10 Metalloids • Found on the stair-step line between metals and nonmetals. • Metalloids have properties that can be similar to both metals and nonmetals – Under some conditions, it can behave as a metal and conduct electricity – Under other conditions, it may not conduct electricity like the nonmetals. • Silicon, a metalloid, is widely used in the computer chip industry. 11 Metalloids 12 Squares in the periodic table • The periodic table displays the symbols and names of the elements, along with the information about the structure of their atoms. 13 Groups or Families • • • • • • Group 1A: Alkali Metals Group 2A: Alkaline Earth Metals Group B: Transition Metals Group 4A – 6A: … Group 7A: Halogens Group 8A: Noble Gases 14 Families (groups) in the periodic Table 15 Electron Configurations in Groups • Electrons play a key role in determining the properties of elements. • There is a connection between the placement on the periodic table and the electron configuration. • We can sort elements based on their electron configuration (See periodic table with electron config) – – – – Noble gases Representative elements Transition Inner transition metals 16 Noble Gases • Found in group 8A; also known as inert gases b/c they rarely take part in a chemical reaction. • The s and p sublevels of the highest occupied energy level are completely filled. – This explains the relative inactivity of the noble gases. Helium Neon Argon Krypton 1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p63d104s24p6 17 The representative elements • Groups 1A, 2A……3A,4A,5A,6A,7A – Does not include B middle and noble gases • Why are they the representative elements? – They display a wide range of physical and chemical properties – Some are metals, some are metalloids, and some are nonmetals – Most are solids, few are gases, and 1 is a liquid. • The s and p sublevels of the highest occupied energy level are not filled. – They participate in reactions when they are able to lose or gain electrons to have full outer energy levels. 18 The representative elements Element Electron Configuration Lithium 1s22s1 Sodium 1s22s22p63s1 Potassium 1s22s22p63s23p64s1 Carbon 1s22s22p2 Silicon 1s22s22p63s23p2 Germanium 1s22s22p63s23p63d104s24p2 19 Transition and Inner Transition Elements • Transition Metals – Group B displayed in the middle of the main body of the periodic table – S sublevels of the higher energy level and a nearby d sublevel contain electrons • Inner Transition Metals – Appear below the main body of the periodic table – S sublevels of the higher energy level and a nearby f sublevel contain electrons Element Electron Configuration Titanium 1s22s22p63s23p64s23d2 (transition) Cerium 1s22s22p63s23p64s23d104p65s24d105p66s24f2 (inner transition) 20 Blocks of Elements • The periodic table is divided into patterns of electron configurations – – – – s block = groups 1A and 2A (including Helium) p block = groups 3A, 4A, 5A, 6A, 7A, 8A d block = transition metals f block = inner transition metals 21 Blocks in the periodic table 22 Abbreviated electron configurations 23 Trends in atomic size • Atomic radius – ½ the distance between the nuclei of two atoms of the same element when the atoms are joined. • 1 meter = 1012 picometers • In general, atomic size tends to – Increase from top to bottom within a group – Decrease from left to right across a group 24 Trends in atomic size 25 Trends in Atomic Size • Why does it increase down a group? – Down a group you have 2 opposite effects: • Nuclear attraction: more protons more positive charge; nucleus pulls electrons in • Shielding effect: increased # of energy levels shields the electrons from the nucleus. – The shielding effect is larger than the nuclear attraction. The atom size gets bigger • Why does it decrease across a period? – Across a group 2 things are occurring: • Nuclear attraction: more protons more positive charge; nucleus pulls electrons in • Shielding effect: CONSTANT… All electrons are in the same energy level – The nuclear attraction is stronger and pulls the electrons in making the atom smaller. 26 Ions • An ION is an atom or group of atoms that has a positive or negative charge. • Positive and negative ions form when electrons are transferred between atoms. – Cation: ion with a positive charge – Anion: ion with a negative charge • Metals become cations by losing electrons • Nonmetals become anions by gaining electrons. 27 Trends in ionic size • Cations are smaller than the atoms from which they form. – When an atom loses an electron, the remaining electrons are drawn closer to the nucleus (increased attraction). • Anions are always larger than the atoms from which they form. – When an atom gains an electron, the attraction of the nucleus for any one electron decreases allowing the size to get bigger. 28 Trends in Ionic Size 29 Trends in ionization energy • If there is enough energy, an electron can be removed from an atom. • Ionization energy: – The energy required to remove an electron from an atom. • First ionization energy tends – to decrease from top to bottom within a group – To increases from left to right across a group 30 Ionization Energy Trend 31 Trends in Ionization energy • Why does it decrease down a group? – It is easier to remove electrons because the shielding effect is stronger than the nuclear attraction hence less energy is required. • Why does it increase across a period? – Across a period, the shielding effect remains constant but the nuclear attraction increases. It takes more energy to remove an electron from the atom. 32 Trends in electronegativity • Electronegativity: The ability of an atom of an element to attract electrons when the atom is in a compound. – Electronegativity can help predict the type of bond (ionic or covalent) that will form during a reaction. – The noble gases do not form many compounds and thus they do not have electronegativity values. • In general, electronegativity values tend to – Decrease from top to bottom within a group – Increase from left to right across a period 33 Trends in electronegativity 34 Trends in electronegativity 35 Summary of Trends 36