Chapter 2 Study Guide

chemistry

has been important since ancient times

• the processing of natural ores to produce metals for ornaments and weapons and the use of embalming fluids are just 2 applications of chemical phenomena that were utilized prior to 1000 BC.

Greeks

• were the first to try to explain why chemical changes occur

• 400 BC: proposed that all matter was composed of 4 fundamental substances: fire, earth, water, and air

• considered the question of if matter was continuous, thus infinitely divisible into smaller pieces or composed of small indivisible particles

• Demokritos and Leucippos, who used the term atom to describe these particles

  • however didn’t have the ability to test their theory

alchemy

the next 2000 years of chemical history were dominated by a pseudoscience called

• some alchemists were mystics/fakes who wanted to turn cheap metal into gold

• many were very serious scientists and saw a period of important advances

• they discovered several elements and learned to prepare the mineral acids

The foundations of chemistry were laid in the 16th century

• With the development of systematic metallurgy (extraction of metals from ores) by a German, Georg Bauer, and the medical application of minerals by a Swiss alchemist/physician know as Paracelsus

Robert Boyle

• The 1st chemist to perform truly quantitative experiments

• He carefully measured the relationship between the pressure and volume of air

• Published the book: The Skeptical Chymist. The quantitative sciences of physics and chemistry were born

  • focused on quantitative behavior of gases and ideas of chemical elements

  • a substance was an element if it could be broken down (killed G idea)

18th century

• combustion studied

• more elements found (basics)

• Antoine Lavoisier (French), explained the true nature of combustion

  • mass was never created nor destroyed

law of conservation of mass

• founded by Lavoisier

• Mass was never created nor destroyed in a chemical reaction

Lavoisier

• his quantitative experiments

Proust

• showed that a given compound always contains exactly the same proportion of elements by mass

law of definite proportion

A given compound always contains exactly the same proportion of elements by mass

Dalton

• elements were composed of tiny individual particles, a given compound should always contain the same combination of these atoms

  • why the same relative masses of elements were always found in a given compound

• used atomic masses

law of multiple proportions

when 2 elements form a series of compounds, the ratios of the masses of the second element that combine with 1g of the 1st element can always be reduced to small whole numbers.

example: 1.750g,0.8750,0.4375g

1.750/0.8750=2/1=2, 0.8750/0.4375=2/1=2,1.750/0.4375=4/1=4

Dalton’s Atomic Theory

1. Each element is made up of tiny particles called atoms

2. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways

3. chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms'

4. chemical reactions involve the reorganization of the atoms-changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction

Gay-Lussac

• Preformed experiments in which he measured (same conditions for T and P) the volume of gases that reacted with each other

Avogadro

• used the results to propose that at the same temperature and pressure, equal volumes of different gases contain the same number of particles called Avogadro’s hypothesis

• this only makes sense if the distance between the particles in a gas are very great compared with the sizes of the particles (gas=depend on the # of molecules present, not their size)

If Avogadro’s hypothesis is correct, Gay-Lussac’s result:

2vol of hydrogen react with 1vol of oxygen —> 2vol of water vapor can be expressed as follows: 2 molecules of hydrogen react with 1 molecule of oxygen —> 2 molecules of water

• these observations can be explained by assuming that gaseous hydrogen, oxygen, and chlorine are all composed of diatomic molecules: H₂

• also found H₂O

19th century

• list of relative atomic masses could be determined

• Berzelius: discovered the elements cerium, selenium, silicon, and thorium and developed the modern symbols for the elements used to write formulas on compounds

electron

• the first important experiments that lead to the understanding of the atom were done by JJ Thomson

Thomson

• who studied electrical discharges in partially evacuated tubes called cathode-ray tubes

• found that when high voltage was applied to the tube, a “ray” or cathode ray was produced

• b/c this ray was produced at the negative electrode and was repelled by the negative pole of an applied electric field

• result: ray was stream of negatively charged particles: electrons

measured the deflection of the beam of electrons in a magnetic field: charge-to-mass ratio of an electron:

e/m=-1.76 × 10^8 C/g

• e=the charge on the electron in coulombs (C)

• m=the electron mass in g

Thomson’s goal

• understand the structure of the atom

• reasoned that since electrons could be produced from electrodes made of various types of metals, all atoms must contain electrons

• atoms=electrically neutral, so they must have a positive charge

• plum-pudding model (cloud + and random spots -)

Milikan

• performed experiments using charged oil drops, allowed him to determine the magnitude of the electron charge

• mass of an electron= 9.11 × 10³¹kg

late 19th century

scientists discovered that certain elements produce high-energy radiation

Bacquerel

• found uranium could produce its image on a photographic plate in the absence of light

• spontaneous emission of radiation of uranium, called radioactivity

early 20th century studies found 3 types of radioactive emissions

gamma rays (y)-high energy light

beta particles (b)-high-speed electron

alpha particles (a)- 2+ charge ( of - (x2) and +) mass: 7300 times that of an electron

Rutherford

• Performed many of the pioneering experiments to explore radioactivity; carried out an experiment to test Thomson’s model.

• directing a particles at a thin foil to crash like cannonballs through the gauze, he expected the alpha particles to travel through the foil with, at most, very minor deflections in their paths

Rutherford experiment

• results: most of the alpha particles passed through; many of the particle deflected at large angles, and some were reflected, never hitting the detector.

• plum pudding model wrong

• the lg defections of the a particles could be caused by a center of concentrated positive charge that contains most of the atom’s mass, in mostly open space, had a closed encounter with a positive center, reflected= direct contact

• found that there had to be a nuclear atom- an atom with a dense center of positive charge (the nucleus) with electrons moving around it at a distance that is relative to the nuclear radius (10^-8)

chem of an atoms results from

its electrons -1

• fine with a crude nuclear model

• nucleus= protons =1 (+ charge equal to the magnitude to the electron’s - charge) and neutrons (same mass as a proton but no charge)

nucleus

1. small size compared with the overall size of the atom

   1. extremely high density

atoms have different chemical properties b/c

the number and the arrangement of electrons

• e- constitute most of the atomic volume and thus are the parts that intermingle when atoms combine to form molecules

• the number of electrons possessed by a given atom greatly affects its ability to interact with other atoms

• diff atoms= diff # of protons and electrons= diff chemical behavior

the protons

equals the number of electrons as atoms have no net charge

• diff types of sodium atoms have diff numbers of neutrons

isotopes

atoms with the same number of protons but different numbers of neutrons

• symbol: atom is written where the atomic number Z (# of protons) is written as a subscript, and the mass number A ( the total number of protons and neutrons) is written as a superscript, then element symbol to the right of those numbers

  • b/c the chem of an atom is due to its electrons, isotopes show almost identical chemical properties

chemical bonds

the forces that hold atoms together in compounds

covalent bonds

• sharing electrons

• the resulting collection of atoms =molecule

molecules can be represented in many ways

• chemical formula: in which the symbols for the elements are used to indicate the types of atoms present and subscripts are used to indicate the relative numbers of atoms

  • CO₂

• more information about a molecule is give by its structural formula: in which the individual bonds are shown (lines), shape can differ

• dashed line=behind the plane of paper wedge= front

compound composed of molecules

• individual molecules move around as independent units

• structural formula:

  • space filling model: which shows the relative sizes of the atoms as well as their relative orientation in the molecule

  • ball and stick models: are also used to represent molecules

the second type of chemical bond results

• from attractions among ions

• ion: is an atom or group of atoms that has a net positive or negative charge

  • table salt or sodium chloride which forms when neutral chlorine and sodium react

• ions are formed when an electron is transferred from a sodium atom to a chlorine atom (neutrons ignored): with 1e- stripped off, the Na, with its 11 protons and only 10e- now has a net 1+ charge— it has become a positve ion.

  • Na—> Na⁺ + e-

cation

positive ion

if an electron is added to chlorine

18e- produced 1- charge; the Cl has become an ion with a neg charge- an anion (Cl^-)

anions and cations attract each other (opp charges)

• force of attraction between oppositely charged ions= ionic bonding

• STEAL