Transcript Chapter 2
Chapter 2 Atoms and Elements atom- smallest identifiable unit of an element element- a substance that cannot be broken down into simpler substances -there are about 91 naturally occurring elements -scientists have made over 20 elements Modern Atomic Theory and Laws That Led to It 1) Law of conservation of mass- matter is neither created nor destroyed -when a reaction is carried out, the mass of the reactants must equal the mass of the products ex- 2Na + Cℓ2 2NaCℓ 7.7g 11.9g 19.6g 19.6g 2) Law of Definite Proportions- all samples of a given compound have the same proportions of their constituent elements ex- 18.0g H20 results in 16.0g O2 and 2.0g H2 mass ratio = 16.0g O2/2.0g H2 = 8 O2:1 H2 -this is true for all samples of water Page 49 Ex 2.1 and For practice 2.1 3) Law of Multiple Proportions- when two elements (A and B) form two different compounds, the mass of element B that combines with 1g of element A can be expressed as a ratio of small whole numbers Page 50 Ex- 2.2 For practice 2.2 John Dalton’s Atomic Theory 1. All elements are composed of atoms 2. All atoms of the same element have the same properties and are identical 3. Atoms combine in simple whole number ratios with other atoms to form compounds 4. Atoms of one element cannot be changed into atoms of another element -Scientists began to think that atoms were composed of smaller particles Discovery of Electron -J.J. Thomson (1856-1940) -used a cathode ray tube -believed that the cathode ray was composed of tiny particles with an electrical charge -these are electrons -rays travel from – charged electrode (cathode) to + electrode (anode) Discovery of Charge of Electron -Robert Millikan (1868-1953) -said electrons have a negative charge Plum Pudding Model -it was known that – charged particles attract + charged particles -it was also known that atoms were neutral -so there must be a + charged particle in the atom -Thomson proposed that – charged electrons were small particles held within a + charged sphere -protons discovered by Eugen Goldstein (1850 – 1930) Ernest Rutherford (1871-1937) -performed experiments that allowed him to conclude that the atom must have a positive mass in a much smaller space than proposed Nuclear Theory 1) most of atom’s mass and all of its + charge are in a small core called the nucleus 2) most of volume is empty space through which negative particles move 3) there are as many negatively charged particles as there are positive (protons) so the atom is neutral -There were still some parts of atom missing James Chadwick (1891-1974) -discovered that missing mass was neutrons Subatomic Particles and the Atom protons (p+) – positively charged, found in nucleus, mass = 1.67 x 10-24g neutrons (n0) – neutrally charged, no charge, found in the nucleus, mass = 1.67 x 10-24g electrons (e-) – negatively charged, found surrounding nucleus in clouds or energy levels, mass = 9.11 x 10-28g **most of mass in the nucleus of an atom atomic number- # of protons in an element, defines the element, the smaller # on the periodic table -since atoms are neutral: # of p+ = # of e-so the atomic # also tells you the # of eHow many protons in the following? argon uranium iron lithium 18 92 26 3 *each element also has that # of e- mass number- sum of the # of p+ and n0, larger # on the periodic table rounded to a whole # How do you find the # of n0? # n0 = mass # - atomic # How many neutrons in the following? argon uranium iron lithium 22 146 30 4 Can be written this way: Ar-40 U-238 Fe-56 Li-7 atomic mass- larger # on the table not rounded, average mass of all the isotopes of that element isotopes- atoms with the same # of p+ and e- but different # of n0 exNe-20 Ne-21 Ne-22 p+ = 10 p+ = 10 p+ = 10 e- = 10 e- = 10 e- = 10 n0 = 10 n0 = 11 n0 = 12 Page 59 Ex 2.3 and For Practice 2.3 natural abundance- the relative % of an isotope with respect to other isotopes of the same element -back to Ne isotope Which isotope of Ne would have the highest relative abundance? Ne-20, because it is closest to the atomic mass -Relative Abundance problems Modern Periodic Table -grew from the work of Dmitri Mendeleev -arranged elements according to increasing mass -called periodic law- when elements are arranged in order of increasing mass/atomic #, similar properties will recur periods- horizontal rows groups/families- vertical columns, indentified by a # and a letter Groups of the Periodic Table Alkali Metals- Group 1A elements Alkaline Earth Metals- Group 2A elements Transition Metals- Group B elements, bridge Inner Transition Metals- two rows at the bottom non-metals- upper and lower right side Metalloids- staircase- starting at boron Halogens- Group 7A Noble Gases- Group 8A or Group 0 -also called inert gases- they do not react to form compounds Main group or Representative elements -Groups 1A to 8A minus the Transition metals Properties of Metals 1) conduct electricity 2) ductile (can be drawn into wires) 3) shiny/lustrous 4) malleable (can be pounded into sheets) -Most elements are solids Gases = H, O, N, F, Cℓ, He Ne, Ar, Kr, Xe, Rn Liquids = Hg, Br (Ga, Fr, Cs)