1.2 Atoms Isotopes and Ions PDF

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Notes about the isotopic symbols used to describe the the different sizes of atoms found in nature....

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Name: ______________________ Chemistry 1 Notes

Topic 1: The Atom 1.2—Atoms, Isotopes, and Ions I.

Atomic Arrangement

 Within the nucleus are protons and neutrons:

proton neutron

mass 1 1

electron 1/2000

 Protons (p+) have a positive charge equal to the negative charge of an electron. Protons are about 2000x heavier than electrons.  Neutrons are electrically neutral particles with effectively the same mass as protons.  Electrons (e–) are negatively-charged subatomic particles that orbit the nucleus of an atom in “clouds”. We need to discuss this electron cloud in terms of probability model of where the electron(s) are most likely to be found at any given point in time…

 atomic number (Z) – the number of protons in the nucleus of each atom of that element  # of protons = # of electrons (since all atoms are electrically neutral)  mass number (A) – the total number of protons + neutrons in the nucleus of an isotope  # of neutrons = mass # – atomic #  Isotopes are atoms of the same element that have different masses.  Isotopes have different masses due to different #s of neutrons. There are three isotopes of hydrogen: protium (99.985%), deuterium (0.015%), and tritium (very rare, radioactive)

charge +1 0 –1

Name: ______________________ Chemistry 1 Notes  2 ways to name isotopes:

A mass# 1) nuclear symbol/isotopic symbol: atomic#symbol = Z ex:

E

238 92 U

2) hyphen notation: element–mass # ex: uranium–238 or U–238 carbon–12 or C–12  So, the three isotopes of hydrogen could be named:  protium: 11H or hydrogen–1  deuterium:  tritium:

3 1H

2 1H

or hydrogen–2

or hydrogen–3

 atomic mass unit (u) – 1/12 the mass of a carbon-12 atom  relative atomic mass – mass of an atom expressed in atomic mass units (it’s relative to the mass of carbon-12)  Non-integer atomic masses (average atomic masses) result from the existence of several different isotopes of that element.  average atomic mass – the weighted average of the masses of the naturally occurring isotopes of an element  For example, there are 2 main isotopes of chlorine: 75% of all chlorine is

35 Cl , 17

35 17 Cl

&

with a mass of 35u and 25% is

37 17 Cl

37 17 Cl

, with a mass of 37u

So the average atomic mass for chlorine is (.75 x 35u) + (.25 x 37u) = 35.5u  Take carbon as a second example: carbon-12 = 12.0u carbon-13 = 13.003355u

98.90%, in nature 1.10%, in nature

average atomic mass for carbon = (0.9890 x 12.0u) + (0.0110 x 13.003355u) = 12.011 u 

Ions are atoms (or groups of atoms) with a positive or negative charge.  cation – positively charged ion (formed by losing e-)  anion – negatively charged ion (formed by gaining e-) 2

Name: ______________________ Chemistry 1 Notes  valence electrons – electrons in the highest energy level 

Valence e- are the only e- available to be lost, gained, or shared in the formation of chemical compounds.

 As we mentioned, elements usually form ions so as to achieve a full octet of eight valence electrons (thus they have a noble gas electron configuration). Here is a look at how atoms or groups of atoms form ions:  Group 1 metals: +1 ions (ex: Na+) by losing one valence electron → full octet.  Group 2 metals: +2 ions (ex: Ca2+) by losing two valence electrons → full octet.  Group 3 (or 13) metals: +3 ions (ex: Al3+) by losing three valence e – → full octet.  Group 5 (or 15) non-metals: –3 ions (ex: P3–) by gaining three valence e– → full octet.  Group 6 (or 16) non-metals: –2 ions (ex: O2–) by gaining two valence e – → full octet.  Group 7 (or 17) non-metals: –1 ions (ex: F–) by gaining one valence e – → full octet.  The transition metals can often form more than one ion (multiple oxidation states): iron forms Fe +2 and Fe+3 ions, copper forms Cu+ and Cu2+ ions, etc. 

Electron-dot notation  Shows only the electrons occupying the highest (or outermost) main energy level. bromine =  Therefore, there can only be a max of 8 electrons in the highest energy level.  An atom with a full 8 electrons has a full octet.  All atoms would like to have a full octet, because that is the most stable arrangement.

3

Name: ______________________ Chemistry 1 Notes

II.

Metallic Bonds

 Metals readily give up electrons (having low ionization energies—meaning it does not take a lot of energy for them to give up electrons), but don’t attract electrons as strongly as non-metals (which have high electronegativities—meaning they cannot attract electrons very well).  In a pure chunk of metal, the valence electrons are delocalized – they do not belong to any one single atom, but are free to move around  metal cations and free electrons.  All these mobile, delocalized valence electrons are shared equally by all the atoms, in a sort of “electron sea”. Thus, a metallic bond results from the attraction between positive metal ions and the surrounding sea of mobile electrons.

 The higher the + charge of the metal cation, the stronger the metallic bond formed. (ex: Al atoms forms stronger metallic bonds than Na atoms, b/c Al forms +3 ions, while Na forms +1 ions. The greater strength of the Al metal bond is due in part to a greater # of delocalized valence e-.)  Consider the properties of metals, and how their “electron sea” helps explain some of them:  high electrical conductivity – if the delocalized electrons are free to move around, that allows a current of electricity to flow easily  thermal conductivity – metals conduct heat as well  When a metal bar is heated at one end, the electrons at that end increase in kinetic energy and rapidly flow through the metal, transporting heat energy with them.  (By contrast, molecular or ionic compounds can only pass along heat slowly, from one neighboring atom to the next.)  malleability – metals can be shaped or hammered into a flat sheet  Again, because of the delocalized electrons, metallic bonding is not directional, but uniform throughout the solid. One plane of metal ions can slide past one another without encountering any resistance or breaking any bonds.

 ductility – metals can be drawn out into a fine wire  This property is explained by the same uniform bonding of metal ions mentioned under malleability.  high melting & boiling point – The metallic bond is an extremely strong bond – stronger than most covalent and ionic bonds. Thus, much energy is required to overcome the metallic bonding, resulting in high melting and boiling points for most metals.  luster – Luster refers to the “shininess” of metals. The delocalized electrons in metals do really absorb or capture light. To the extent that they do absorb light, the electrons quickly return to their original energy level by emitting the energy they absorbed with little dissipation. So, when light strikes a metal surface all visible wavelengths are reflected with little or no alteration in color. 4...