Chapter 4: Atoms and Elements

Atoms

the smallest identifiable unit of an element

• compose matter

• their properties determine the properties of matter

• 91 kinds

• spherical shape

element

a substance that cannot be broken down into simpler substances

• 91 found in nature

• 20 synthetic

• exact number controversial because some previously considered only synthetic may actually occur in nature in very small quantities

• defined by their number of PROTONS in the nucleus of an ATOM

  • if the number of protons change it would be a different element

Democritus Atomic Theory

• first person to postulate that matter is composed of atoms

• him and his mentor Leucippus recorded ideas of atoms

• suggested that if you divide matter into smaller and smaller pieces you end up with tiny indestructible particles

• atomos: invisible

Dalton Atomic Theory

• formalized this theory 2000 years after Democritus

1. Each element is composed of tiny, indestructible particles called atoms

2. all atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements

3. atoms combine in simple whole number ratios to form compounds

J.J Thomson

discovered a smaller and more fundamental particle called the electron

• discovered:

  • electrons are negatively charged

  • electrons are much smaller and lighter than atoms

  • electrons are uniformly present in many different kinds of substances

  • proposed that atoms must contain positive charge that balances the negative charge of electrons

Plum pudding model of the atom

in the model suggested by JJ Thomson, negatively charged electrons were held in a sphere of positive charge

Rutherford’s gold foil experiment

tiny particles called alpha particles were directed at a thin sheet of gold foil

• most of the particles passed directly through the foil

• a few were deflected—some of them at sharp angels

Discovery of the atomic nucleus

1. expected result of Rutherford’s gold foil experiment

   1. if the plum pudding model correct the alpha particles would pass right through the gold foil with minimal deflection

2. actual result of Rutherford’s gold foil experiment

   1. a smaller number of alpha particles were deflected or bounced back

Rutherford Nuclear Theory of the Atom

1. most of the atoms mass and all of its positive charge are contained in a small core called the nucleus

2. most the the volume of the atom is empty space through which the tiny, negatively charged electron are dispersed

3. the number of negatively charged electrons outside the nucleus is equal to the number of positively charged particles (protons) inside the nucleus so that the atom is electrically neutral

Distribution of mass in the atom

• the nucleus makes up 99.9% of the atoms mass and occupies a small fraction of its volume

• electrons are distributed through a much larger region but don’t have much mass

• matter at its core is less uniform that it appears

  • if matter were composed of atomic nuclei piled on top of each other like marbled, it would be incredibly dense

  • a single grain of sand composed of a solid atomic nuclei would have a mass of 5 million kg

  • astronomers believe that black holes and neutron stars are composed of this kind of incredibly dense matter

electrical charge

a fundamental property of protons and electrons

• positive and negative charges attract each other

• positive-positive and negative-negative charges repel each other

• positive and negative charge cancel each other so that a proton and an electron, when paired, are charge-neutral

Properties of protons, neutrons, and electrons

protons and neutrons: similar masses

electrons have almost negligible mass

atomic number`

the number of protons in the nucleus of an atom; given the symbol Z

Dmitiri Mendeleev and recurring properties

• Russian chemistry professor

• proposed from observation that when the elements are arranged in order of increasing relative mass, certain sets of properties recur periodically

• we arrange the elements in rows so that similar properties align in the same vertical columns

Periodic Law

• Mendeleev’s periodic law based on observation

  • summarized many observations but did not give the underlying reason for the observation—only theories do that

metals

• occupy the left side of the periodic table and have similar properties

• good conductors of heat and electricity

• can be pounded into flat sheets (malleability)

• drawn into wires (ductility)

• shine (lustrous)

• lose electrons when they undergo chemical changes

• ex: iron magnesium, chromium, and sodium

Non metals

• occupy the upper right side of the periodic table

• dividing line between metals and nonmetals is the zigzag line running from boron to astatine

• more varies properties; some are solids at room temperature while other are gases

• as whole, they tend to be poor conductors of heat and electricity

• gain electrons when they undergo chemical changes

• ex: oxygen, nitrogen, chlorine, iodine

Metalloids

• line along the zigzag line dividing metals and non metals

• also called semi metals= mixed properties

• semiconductors—their intermediate electrical conductivity, which can be changed and controlled

  • this makes semiconductors useful in manufacture of electronic devices that are central to computers, cell phones, and other modern gadgets

• ex: silicon, arsenic, and germanium

main group elements

properties can generally be predicted based on the elements position

transition elements

properties tend to be less predictable based on the elements position

ions

in chemical reactions atoms often lose or gain electrons to become charged particles

cations

positive ions (more protons than electrons)

anions

negative ions (more electrons than protons)

making ions by losing electrons

• the charge of an ions depends on how many electrons were gained or lost and is given by the formula

  • ion charge=number of protons-number of electrons

• where p+ stands for proton and e- stands for electron

• ex: in reactions lithium atoms lose one electron (e-) to form Li+ ions

  • 3 protons and 2 electrons ion charge= 3-2= 1+

making ions by gaining electrons

• The charge of an ion depends on how many electrons were gained or lost and is given by the formula:

  • ion charge=number of protons-number of electrons

• fluorine atoms gain one electron (e-) to form F- ions

  • F- ion with 9 protons and 10 electrons ion charge 9-10= 1-

isotopes

atoms with the same number but different numbers of neutrons

• few exceptions to the rule (like boron)

• all atoms of a given element have the same number of PROTONS

• the DON’T have the same number of NEUTRONS

• all elements have their own unique natural abundance of isotopes

• each naturally occurring sample of most elements has the same percent natural abundance of each isotope

• characterized by their mass number (weighted average of the masses of the individual isotopes)

mass number

the sum of the number of protons and the number of neutrons

• number of neutrons is the difference between the mass number and the atomic number

Atomic Theory

• ancient greeks: matter is composed of small indestructible particles

• dalton: matter is composed of atoms

• atoms of a given element have unique properties that distinguish them from atoms of other elements

• atoms combine in simple whole number ratios to form compounds

electrons

negatively charged subatomic particles that make up most of the atoms volume

protons

positively charged subatomic particles that make up the nucleus which is most of the atoms mass

neutron

neutral (no charge) subatomic particle makes up the nucleus which is most of an atoms mass

alkali metals

highly reactive metals (1A)

alkaline earth metals

fairly reactive metals (2A)

Halogens

very reactive nonmetals (7A)

Semiconductor

compound of element exhibiting intermediate electrical conductivity that can be changed and controlled

noble gases

chemically unreactive (8A)

periodic table

• Organized left to right by increasing atomic number (Z)

• Organized up and down by properties (groups/families)

  • More important than atomic number

• Left are metals (tend to lose electrons in chemical changes)

• Upper right are nonmetals (tend to gain electrons)

• Between are metalloids

Robert Millikan

determined the magnitude of the electron’s charge 

• In his experiment, small electrically charged drops of oil were suspended between two metal place 

  • The drops were subjected to the downward force of gravity, and the upward attraction of an electrical field 

  • Once the oil drops move about, Millikan showed that their charge was always a precisely determined charge, the electron’s charge

James Chadwick

discovered the neutron 

• Determined that there is also a neutrally charged subatomic particle in an atom 

  • This particle’s mass was about the same as the proton 

  • Previously assumed this particle was a proton, but he proved himself wrong

Eugen Goldstein

discovered the proton 

• Used a cathode ray tube and noticed that there were rays traveling in an opposite direction from the cathode rays 

  • He called these canal rays and they were composed of positively charged particles 

• Through this experiment, he found out that the atom also included a positively charged subatomic particle, in addition to the negatively charged subatomic particle 

  • Mass of the proton = nearly 2000 times the mass of the electron

blocks of elements

SPDF Block: where the valence electrons are being filled in *refer to picture on study guide

S- 1A, 2A, He 

P- 3A, 4A, 5A, 6A, 7A, 8A (not include He)

D - transition metals

F - last two rows

representative elements

• Main group elements in S and P blocks (except noble gases)

• Completely filled inner orbitals and incomplete outer orbitals

• Ex: Alkali metals (Li), Alkaline earth metals (Be), Halogens (Cl, F)

• 1-8 in roman numerals at the top of each row

• How you know how many valence electrons are available and how you find out the charge 

• Ex: 1A (+1 charge)