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Chapter 6 Periodic trends Mendeleev- organized the periodic table by mass. grouped them by similar properties. - like a game of Solitare. didn’t know about subatomic particles at the time. Exception to his rule: Iodine and tellurium. - Mendeleev thought that he had miscalculated the masses since he knew that iodine belonged in the group with bromine and chlorine. Modern Periodic Table is arranged by atomic number and grouped by properties. Group numbers are standardized by the International Union of Pure and Applied Chemistry (IUPAC) There are 3 Classes of elements. - Metals- conductors, ductile(made into a wire) and malleable( able to resist shattering) - Nonmetals- insulators (except carbon), brittle, mostly gases - Metalloids- have properties of both metals and nonmetals Ex. Silicon- alone it is an insulator (nonmetal) o Combined with boron it is a conductor (metal) Periodic Trends - Group Trends- what happens within the group or column of elements - Period trends – what happens across a period or row. Groups of the periodic table - “A” elements- Representative Elements 1A- Alkali Metals 2A –Alkaline Earth Metals 7A – Halogens 8A – Noble Gases - “B” Elements are the transition metals. - “A” group number corresponds to the number of valence electrons in that group. Valence Electrons- Outer shell electrons only Periodic Trends Atomic radius - Measuring an atom- the distance between two atoms divided by 2 gives the atomic radius. - 3 factors affect atomic radius Energy level- Higher the energy level the further the electrons are from the nucleus. Charge on the nucleus- more valence electrons increase nuclear charge and pull electrons closer to the nucleus. Shielding effect- the ability for electrons to shield the nucleus from other electrons. Group trend – Atomic radius increases going down the group. Each element in the group has another energy level Atom gets bigger with increased energy levels Period trend- Atomic radius gets smaller as elements go from left to right across a period. - increased nuclear charge pulls electrons closer to the nucleus. Ionization energy(IE) - Energy required to remove one electron from a valence shell. Remove 1 e- = 1+ ion - Energy required to remove the 1st electron is the 1st ionization energy - Energy required to remove the 2nd electron is the 2nd ionization energy 2nd IE is always more than the 1st IE. 3rd IE is more than the 2nd IE. - Greater nuclear charge increases IE because atoms that have fuller valences are more reluctant to give up electrons. - Distance from the nucleus decreases IE because electrons in upper energy levels are further away from the nucleus. Group trend- IE decreases Period trend -IE increases Ionic radius – size of an atom once it becomes and ion - Cations (metals) lose electrons when they form ions. - Anions (Nonmetals) gain electrons when they form ions. - As cations lose electrons the ion radius gets smaller. - As anions gain electrons they get bigger. Group trend- increases generally Period trend- Within cations, ionic radius decreased from left to right Within anions, ionic radius decreases from left to right. Electronegativity- the ability of an atom to attract an electron when chemically combined with another element. - Larger electronegativity will pulls electrons toward it. Group trend – electronegativity decreases because electrons are further away from the nucleus. Period trend- electronegativity increases because there are more electrons surrounding the nucleus. Noble gases are not usually included in ionic radius and electronegativity trends because of their lack of reactivity. Summary of trends- pg 178