Chapter 2-1 Atoms, Molecules, & Ions Notes

Atoms, Molecules, & Ions

Early Philosophy of Matter

Differing Views

• Some philosophers, like Democritus (ca. 460 – ca. 370 BCE), proposed that matter consists of ultimate, tiny, indivisible particles called atoms. He stated, “Nothing exists except atoms and empty space; everything else is opinion.”
• Other philosophers, such as Aristotle (385 – 323 BCE), believed that matter was infinitely divisible.
• There was no experimental way to prove either theory correct at the time. The best debater (Aristotle) was assumed correct.
• Atoms weren’t "seen" until the development of scanning tunneling microscopes (STM) and atomic force microscopes in the 1980s. IBM used 35 Xe atoms in the 1980s.

Scientific Understanding of Nature

Philosophical vs. Scientific Approach

• Philosophers aimed to understand the universe through reasoning and thinking about "ideal" behavior, focusing on asking good questions.
• The late 16th century saw the establishment of a scientific approach to understanding nature through observation and experiment.
• Copernicus's "On the Revolutions of the Heavenly Orbs" (1543) placed the sun at the center of the universe, arguing that the Earth moved across the heavens as one of the planets. He waited over 30 years to publish due to the controversial nature of his ideas.

Essence of Science

• Richard Feynman: “It doesn't matter how beautiful your theory is, it doesn't matter how smart you are. If it doesn't agree with experiment, it's wrong.”

Scientific Revolution

Laws

• Law of Definite Proportion – Joseph Proust (1754 – 1826): All samples of a given compound, regardless of their source or preparation, have the same proportions of their constituent elements.
• Law of Multiple Proportions – John Dalton (1766 – 1844): When two elements (A and B) form two different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small, whole numbers.
• Law of Conservation of Mass – Antoine Lavoisier (1743 – 1794): In a chemical reaction, matter is neither created nor destroyed.

Chemistry Defined

• Chemistry is the science that seeks to understand the behavior of matter by studying the behavior of the atoms and molecules it is composed of.

Atoms and Molecules

• Atoms: Fundamental building blocks of all matter and the smallest part of an element.
• Molecules: Two or more atoms attached together by bonds, which have different strengths, shapes, and patterns.
• Serotonin (C10H12N2O): Increased by SSRIs like Celexa or Prozac.

Dalton’s Atomic Theory (1803)

1. Each element is composed of tiny, indestructible particles called atoms.
2. Atoms of an element have the same mass and other properties that distinguish them from atoms of other elements.
3. Atoms combine in simple, whole-number ratios to form molecules.
4. In a chemical reaction, atoms cannot change into atoms of another element but rearrange the way they are attached.

2 H2 (g) + O2 (g) \rightarrow 2 H_2O (l)

Impact

• Unified the Laws of Conservation of Mass, Definite Proportion, and Multiple Proportions.
• Based on matter having ultimate, indivisible particles (atoms).
• Several parts not quite correct.
• Still amazing insight into the understanding of chemistry

Notes on Charge

• Charges: Positive (+) and Negative (–)
• To be neutral, either:
  1. No overall charge
  2. Equal amounts of opposite charges
• Like charges repel (+ repels +, – repels –)
• Opposite charges attract (+ attracted to –)
• Example: magnesium nitrate - Mg(NO3)2

Cathode Ray Tubes

Karl Ferdinand Braun (1850 – 1918)

• German inventor and physicist; Nobel Prize in Physics in 1909
• Contributed significantly to the development of radio and TV technology
• The tubes contain metal electrodes within an evacuated (vacuum) environment.
• When connected to a high voltage power supply, a glowing area is seen emanating from the cathode.

J.J. Thomson (1856 – 1940)

• English physicist; 1906 Nobel Prize
• Investigated the effect of placing an electric field around the cathode ray tube.
• Hypothesis: Light is NOT deflected by an electric field; charged matter is attracted to an electric field.
• Thomson believed that cathode rays are composed of tiny, negatively charged particles (electrons).
• Every material tested contained these same particles.

Charge-to-Mass Ratio of the Electron

J.J. Thomson's Experiment

• Measured the force required to deflect the path of cathode rays.
• The amount of deflection depends on both the charge and mass of the particles.
• Deflection was measured due to both magnetic and electric fields.
• The charge/mass ratio of these cathode ray particles was -1.76 \times 10^8 C/g

Thomson’s Conclusions

• The only way to explain the origin of these charged particles was if they were “broken off–pieces” that came from the atoms.
• The atom is not so unbreakable after all.
• Important Moment:
  • Recognizing when experimental evidence contradicts existing theories.
  • Refining our internal model of how the universe works based on new evidence.
  • This refined model remains subject to further refinement with new contradictory evidence or more accurate measurements.

Revisiting Dalton's Atomic Theory

Dalton’s Atomic Theory (1803):

1. Each element is composed of tiny, but destructible particles called atoms.
2. Atoms of an element have same mass (and other properties) that distinguish them from atoms of other elements.
3. Atoms combine in simple, whole–number ratios to form molecules
4. In a chemical reaction, atoms cannot change into atoms of another element. Rearrange the way they are attached

2 H2 (g) + O2 (g) \rightarrow 2 H_2O (l)

Thompson’s Conclusions

• Thomson believed that these particles (electrons) were therefore the ultimate building blocks of matter!
• Hydrogen ion (a proton) – Charge/mass = m/z = + 9.58\times 10^4 C/g
• Mass of a Hydrogen atom = 1.673 \times10^{-24} g
• If the particle has the same magnitude of charge as a hydrogen ion, then it must have a mass almost 2000x smaller than a hydrogen atom!

A New Hypothesis for the Atom

• Since the atom is no longer indivisible, Thomson proposed a new model: the atom actually has an inner structure.

Thomson’s Plum Pudding Atom

• Assumed no positively charged pieces (none showed up in the cathode ray experiment).
• Atom HAS TO CONTAIN MANY, MANY, MANY negatively charged electrons
• Mass of a Hydrogen atom = 1.673 \times 10^{-24} g
• Mass of an Electron = 9.1 \times 10^{-28} g
• Hydrogen atom must contain ~2000 e–
• Electrons held in the atom by their attraction for a positively charged electric field within the atom
• Source of positive charge because the atom is neutral.
• The atom is mostly empty space

Millikan’s Oil Drop Experiment – Charge of an Electron

• Robert Andrews Millikan (1868 –1953)
• American Experimental physicist, Nobel Prize in physics (1923)

Millikan’s Oil Drop Experiment

• Electrons are particles found in all atoms
• e– charge/mass: -1.76 \times 10^8 C/g (Thompson Experiment)
• e– charge: -1.60 \times 10^{-19} C (Millikan’s Experiment)
• Know the mass/charge ratio AND the charge, calculate the mass of an e–
• mass = 9.1 \times 10^{-28} g

Discovering More Subatomic Particles – Radioactivity

Key Figures

• Henri Becquerel (Nobel Prize Physics 1903) and Marie Curie (Nobel Prizes Physics 1903 & Chemistry 1911)

Radioactivity Defined

• Certain elements constantly emit small, energetic particles and rays.
• These energetic particles could penetrate matter.

Ernest Rutherford (Nobel Prize Chemistry 1908)

• Identified three different kinds of emissions:
  • Alpha (\alpha) particles: mass 4x the hydrogen atom and a positive charge (+2) — (Helium nucleus)
  • Beta (\beta) particles: mass ~1/2000 of that of a hydrogen atom and a negative charge (electron)
  • Gamma (\gamma) rays: high energy rays, not particles
• Alpha (α) Radiation occurs when an unstable nucleus emits an α particle composed of 2 protons and 2 neutrons.

Rutherford’s Experiment

• To show that something is mostly empty space (Thomson’s Plum Pudding Model), Rutherford tried to shoot something through it. If atom was like a plum pudding, all the α particles should go straight through

• Use large target atoms and very thin sheets of target (so do not absorb “bullet”)

  • Gold has a mass of 197 amu & is very malleable (~0.005 mm thick)

• Use very small particle as a bullet with very high energy

  • α particles — mass of 4 amu & charge of +2

Rutherford’s Results

• Ernest Rutherford (1871 – 1937), New Zealand, 1908 Nobel Prize, “The Father of nuclear physics”
• Over 98% of the α particles went straight through
• About 2% of the α particles went through but were deflected by large angles
• About 0.01% of the α particles bounced off the gold foil
• “…as if you fired a 15” cannon shell at a piece of tissue paper and it came back and hit you.”

Nuclear Atom

• Most α particles go straight through
• Some α particles go through, but are deflected
• A few of the α particles do not go through

Structure of the Atom

• Rutherford proposed the nucleus had a particle with the same amount of charge as an electron but opposite sign.
• Nucleus is positively (+) charged.
• Atom is overall electrically neutral (0).
• Amount of positive charge balances the negative charge of the electrons.
• These positively charged sub-atomic particles are called protons
  • charge = + 1.60\times 10^{-19} C
  • mass = 1.67262\times 10^{-24} g
• Since protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons in its elemental form.

Rutherford’s Interpretation – the Nuclear Model

• The atom contains a tiny dense center called the nucleus
  • Typical atomic radius = 100 pm (100\times 10^{-12} m)
  • Radius of nucleus = 5 \times 10^{-3} pm
  • (20,000 times less!)
  • Space taken by the nucleus = 1/10 trillionth (10^{-13}) the volume of the atom
• Since the electron(s) are responsible for almost all the volume of an atom
• Nucleus has essentially the entire mass of the atom

The Nuclear Model

• If an atom was the size of AT&T Stadium, the volume of the atomic nucleus would comparable to that of a large marble.

Relative Charge

• Protons have equal amounts of charge to electron but opposite signs.
• Proton has a charge of +1 charge unit (cu)

There are different subatomic charges in an atom

• The charge of a proton = + 1.60\times 10^{-19} C (Coulombs)
• The charge of an electron = - 1.60\times 10^{-19} C

Relative scale

• Easier to compare things to each other rather than to outside standard
• Compare it to the amount of charge of an electron
• –1 charge unit (cu) By definition

Relative Mass

• We need a relative scale for mass, which isn't as straightforward as charges.
• Atomic Mass Unit (amu) – unit of mass used to express atomic and molecular weights
• A proton has a mass of 1.67262\times 10^{-24} g
• An electron has a mass of 0.00091\times 10^{-24} g

Rutherford’s Atomic Model Addresses Problems

1. How could beryllium (Be) have 4 protons stuck together in the nucleus?
   • Positive (+) repel each other!?!?
2. Beryllium atoms - 4 protons (Each proton weighs ~1 amu), & should weigh ~4 amu; but it actually weighs ~9 amu!
   • Where is the extra mass?
   • Not e- mass is only 0.00055 amu
   • 4 electrons can’t account for the extra 5 amu of mass

Solution: Neutrons

• There Must Be Another Subatomic Particle
• Rutherford proposed that there was another particle in the nucleus Called a neutron
• 0 charge and ~ 1 amu mass
• Mass = 1.67493\times 10^{-24} g (Slightly heavier than a proton)

Subatomic Particles

Carbon-12

• Carbon–12 (C–12) has 6 protons, 6 neutrons, and 6 electrons
• 1 atomic mass unit (1 amu) is defined as 1/12th the mass of a C–12 atom
• Therefore a C–12 atom = EXACTLY 12 amu (by definition)
• Mass of C–12 = ΣMass Protons + ΣMass Neutrons + ΣMass of Electrons
• Mass of C–12 = (6)(1.00727 amu) + (6)(1.00866 amu) + (6)(0.00055 amu) =12.09888 amu
• But- Mass of C–12 = 12 amu (by definition) not 12.09888 amu
  Where is the “lost mass” of C–12?

A Second Issue With Dalton’s Atomic Theory

• To explain these laws, the matter is based on ultimate, indivisible particles

Several parts of Dalton’s atomic theory are not quite correct (noticing a pattern here)

• Thompson: Cathode Ray
• Rutherford: Gold Foil, neutron
• Curie and Others: Nuclear Chemistry

Elements

• The Number of Protons Defines the Element
• Each element has a unique number of protons in its nucleus The number of protons in the nucleus of an atom is called the atomic number
• Each element has a unique atomic number, name, and symbol
• The elements are arranged on the Periodic Table in order of their atomic numbers

Isotopes

• Frederick Soddy (1877 – 1956) discovered in 1913 that the same element can have different masses, these are called isotopes.

• Example: 35Cl and 37Cl, one with a mass of ~35 amu and another with a mass of ~37 amu.

• The observed mass is a weighted average of the masses of all the naturally occurring atoms, it is called the %percentage of natural abundance of an isotope.

• The atomic mass of chlorine is 35.45 amu

• All isotopes of an element are chemically identical, undergo the same chemical reactions, but they can be separated.

• All isotopes of a specific atom, in its elemental form, have the same number of protons and electrons (e.g. chlorine; 17 p+ and 17 e– ). isotopes have different masses => different isotopes have different numbers of neutrons

• Isotopes are identified by their mass numbers (# protons + # neutrons)

  • Atomic Number (Z) — Number of protons

  • Mass Number (A) — Number of protons + neutrons

  • Cl–35

  • Cl–37 There are 2 isotopes of chlorine found in nature, One mass of about 35 amu Another mass about 37 amu, 3517Cl 3717Cl

Some Important Isotopes and Their Uses

• Hydrogen has three isotopes:
  • Protium (H–1): the most common isotope of hydrogen, only 1 proton
  • Deuterium, D, (H–2): It has 1 proton and 1 neutron. Used as a marker to study the course of chemical reactions and is used in heavy water D2O, which is used in nuclear reactors.
  • Tritium, T, (H–3): has 1 proton and 2 neutrons. Unlike protium and deuterium, tritium is radioactive.
• Uranium consists of two isotopes, U–238 and U–235. Uranium is enriched to increase the content of U–235 which undergoes fission more easily for use in nuclear reactors or to make nuclear weapons
• Carbon has 2 stable isotopes, C–12 (~99%) and C–13 (~1%) and a much rarer radioactive isotope, C–14. The radioactivity of C–14 makes carbon dating possible.

Practice Question:

• How many protons, electrons, and neutrons are in an atom of Cr (Cr–52 )? Atomic number = 24
• Atomic mass = 52
• Mass = m = p+ + n0
• n0 = m - p+
• n0 = 52 – 24
• n0 = 28