Chapter 4 Atoms and Elements PDF

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2/20/21 Chem 101 Chapter 4: Atoms and Elements 

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Dalton’s Atomic Theory o Dalton’s Model  Elements are indivisible, indestructible pure matter, atoms are the smallest pieces of elements (units of elements)  Atoms of a given element are identical (having the same size, mass, and chemical properties), different elements are made from atoms with different properties  Compounds are made from atoms of more than one element that are combined  Compounds contain atoms in small whole-number ratios  Atoms can combine in different ratios to make different compounds  Dalton’s atoms were individual particles  Atoms of each element are alike in mass and size  Atoms of different elements are different in mass and size  Dalton’s atoms combine in specific ratios to form compounds  Two elements of Dalton’s atoms can combine in different ratios to form different compounds o The Law of Multiple Proportions  Atoms of two or more elements may combine in different ratios to produce more than one compound Elements and the Periodic Table o There are 118 elements (types of atoms) that are known to exist  The periodic table lists all the known elements, one per box  Each element is given a name and an abbreviation known as a chemical symbol  Some of them are not found naturally  Chemists make them and they exist for tiny fractions of one second  Committees decide if they are real and name them o Chemical symbol rules  All symbols have either one or two letters  If one letter is used, it is capitalized  If two letters are used, only the first letter is capitalized  Some chemical symbols are east to guess, and others appear to make no sense (come from Latin name abbrev.) o Elements needed to memorize  H Hydrogen 1  He Helium 2  Li Lithium 3  Be Beryllium 4  B Boron 5

 C Carbon 6  N Nitrogen 7  O Oxygen 8  F Fluorine 9  Ne Neon 10  Na Sodium 11  Mg Magnesium 12  Al Aluminum 13  Si Silicon 14  P Phosphorus 15  S Sulfur 16  Cl Chlorine 17  Ar Argon 18  K Potassium 19  Ca Calcium 20  Ti Titanium 21  V Vanadium 23  Cr Chromium 24  Mn Manganese 25  Fe Iron 26  Co Cobalt 27  Ni Nickel 28  Cu Copper 29  Zn Zinc 30  As Arsenic 33  Br Bromine 35  Kr Krypton 36  Rb Rubidium 37  Sr Strontium 38  Ag Silver 47  Sn Tin 50  I Iodine 53  Xe Xenon 54  Cs Cesium 55  Ba Barium 56  W Tungsten 74  Pt Platinum 78  Au Gold 79  Hg Mercury 80  Pb Lead 82  Rn Radon 86  Fr Francium 87 o The Periodic Table  Mendeleev and Mosley are credited with the early periodic tables  Many scientists tried to arrange the elements into some kind of order

Vertical columns are called groups  Elements in the same periodic group behave very similarly  Four periodic groups have special names  Rows are called periods  The elements on the very bottom rows are actually part of periods 6 and 7, but cannot fit easily  The portions of those periods that are removed are given special names  From left to right, and top to bottom, the elements are arranged in order by a whole number called an atomic number  As the atomic number increases, often the atomic mass of the elements also increases, but there are many exceptions  The staircase divides the table into two types of elements (metals and nonmetals) o Groups and series names to know  Group 1 = alkali metals  Group 2 = alkaline earth metal  Group 17 = halogens  Group 18 = noble gases  First row on bottom = lanthanide series  Second row on bottom = actinide series  Elements left of the staircase, except H = metals  Elements touching the staircase, except Al = metalloids  Elements right of the staircase = nonmetals and H o Metals, nonmetals, and metalloids  While defined solely on the elements’ position in the periodic table, metals share some common properties as do nonmetals  Metals o Defined as being to the left of the staircase o Tend to form cations (lose electrons) and ionic bonds o Reacts with nonmetals o Usually, solid o Has a luster (metallic shine) o High density o High melting point o Conducts electricity and speech o Malleable (pliable)  Nonmetals o Defined as being to the right of the staircase line o Can form anions (gain electrons) through ionic bonds can also form covalent bonds (share electrons) o React with nonmetals and metals o Can be solid or gas o Dull appearance o Low density 

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o Low melting point o Nonconductive o Brittle Metalloids o Defined by as touching the staircase line and such as like both metals and nonmetals  Ones on the right act more like nonmetals and ones on the left act more like metals

Atomic Models o Periodic table is more than just a list of elements  There are lots of patterns in the periodic table which is very helpful  To understand the patterns, we have to first learn the structure of atoms o Models of the Atom tutorial on moodle  Dalton’s model of the atom (Billiard Ball model)  Concepts of the atom and compounds  Thomson’s model of the atom (Plum Pudding model)  Electrons in atom  Rutherford’s model of the atom (nucleus model)  Positively charged nucleus core  Bohr’s model of the atom  Electrons occupy orbitals o Deflection and scattering of alpha particles by positive nuclei  Deflection- repelled back to where it started  Scattering- still pretty straight path, but steered off course o Conclusion of atomic models  What are atoms made of?  Evidence from the late 1800s showed that atoms were made of subatomic particles (atoms were in fact divisible) o Faraday and Arrhenius discover ions (charged particles) o 1897 Thomson discovered atoms contain negatively charged particles (electrons) o 1910 Rutherford discovered most of the atom is empty space; later in 1910 Rutherford discovered atoms contain positively charged particles (protons) o Rutherford also found the protons were much heavier than electrons are were located in a core of mass in center of atom (nucleus) o 1932 Chadwick discovered uncharged particles in the nucleus (neutrons), they were much heavier than electrons o Atoms are made of other particles  The reason an atom is uncharged overall is because the electron charge cancels out the proton charge

An atom contains the same number of electrons as protons  Proton- + charged particle found in the core (nucleus)  Neutron- uncharged particle found in nucleus  Electron- - charged particle found in outer area  Since neutrons are uncharged, an atom can have a different number of neutrons then protons  The number of protons in an atom is its atomic number  Atoms with different numbers of protons will be different elements o Summary of subatomic particles 

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Subatomic Particle

Mass

Charge

Located

What happens to an atom if lost or gained?

Electron (e-)

~ 0 amu (1/1780 amu) 9.10 x 10^-24 g

-1 -1.602 x 10^-19 C

Orbitals (far from nucleus)

Proton (p+) 1 amu +1 1.672 x 10^-24 g -1.602 x 10^-19 C

Nucleus

Neutron (n0)

Nucleus

-adding e- makes anion -losing e- makes cation -can’t gain or lose p+ -different elements are atoms with different # of protons -can’t gain or lose n) -isotopes are atoms (of the same element) with different # of neutrons

1 amu 0 1.647 x 10^-24 g

Isotope Symbols o Example:  Lithium  Atomic number of 3 means 3 protons  Mass of 7 amu o 3 protons, each with a mass of 1 amu, plus 4 neutrons, also with a mass of 1 amu each, which adds up to 7 amu  Electron masses are too tiny to be relevant  Boron  Atomic number of 5 means 5 protons  Mass of 9 amu

o 5 protons, each with a mass of 1 amu, plus 4 neutrons, also with mass of 1 amu each, which adds up to 9 amu o Isotopes  Two atoms of the same element, with different numbers of neutrons, and thus with different mass numbers  Isotopes are found naturally  Under normal conditions, you cannot add neutrons to or subtract neutrons from an atom to make an isotope o Exception is nuclear reactions  Symbols  Mass number = # of protons + # of neutrons o Also known as “A”  Note: mass # is equal to that one atom’s mass in amu  Atomic number = # of protons in an atom o Represented by “Z”  Note: atomic number uniquely ID’s an element  Isotopes- two atoms of the same element with different number of neutrons and different mass numbers o Ex. 23Na and 22Na  Mass number minus atomic number equals number of neutrons in the nucleus  Mass number (# of protons + # of neutrons) – atomic number (# of protons) = number of neutrons o Ex. Ag  Mass number 109  Atomic number 47  109 – 47 = 62 neutrons  Converting isotope symbols to numbers of protons, neutrons, and electrons  Ex. 13C o # p = atomic number (from table) = 6 o #n = mass number – atomic number = 13 – 6 = 7 o #e = same as # p if its an atom (uncharged) = 6  Atoms can form ions (charged particles) by gaining or losing electrons  When one or more electrons are lost from an atom, a cation is formed  When one or more electrons are gained by a neutral atom, an anion is formed o Remember… protons can never be gained or lost

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Average Atomic Mass o Math review…  It is possible to have an average that is not equal to any of the numbers that were averaged:  76, 89, 96, 81 o Average is 85.5 or 86 o Average atomic mass  Although there is no actual Cl atom with a mass of 35.5 amu, the average mass of one Cl atom in a sample with millions of Cl atoms is 35.5 amu o Make visual of large sample of atoms 75% Cl-35  Sample of 4 atoms o 35Cl, 35Cl, 35Cl, 37Cl o Average = (35+35+35+37)/4 = 35/5 amu  Sample of 12,364,287 atoms (too many to count) o 75% are 35Cl and 25% are 37Cl o Average = (.75 x 35) = (,25 x 37) = 35.5 amu  The average atomic mass has been calculated for all elements based on the natural abundance of each isotope and its isotope mass and it appears on the periodic table (non-whole number) o Calculating average atomic mass  Need to know natural abundance % and actual isotope mass (not exactly equal to the mass number) for each isotope  Ex. Silicon o Mass number 28, 29, 30 o Isotope mass 27.976927, 28.976495, 29.973770 o % abundance 92.23, 4.67, 3.10  For each isotope, multiply the decimal of the abundance by the isotope mass  Then add results  (0.9223 x 27.976927) + (0.0467 x 28.976495) + (0.0310 x 29.973770) = 25.80 + 1.35 + 0.0929 = 28.08...