Atoms/molecules have a lot of space between them and move freely.
Highly compressible.
2. Classification of Matter by Components
Matter can be classified as:
pure substance vs mixture.
If a pure substance, one type of particle.
If a mixture, two or more components in variable proportions.
Pure Substances vs Mixtures
Pure Substance
Invariant composition; made of only one component.
Mixture
Composition can vary from sample to sample; made of two or more components.
Classification of Pure Substances
Pure substances split into two types:
1) Element: cannot be chemically broken down into simpler substances.
2) Compound: composed of two or more elements in fixed definite proportions.
Most elements are chemically reactive and form compounds (e.g., water, sugar).
Classification of Mixtures
Heterogeneous mixture: composition varies by region (not uniform). Examples include wet sand.
Homogeneous mixture: appears one substance; uniform composition; atoms or molecules mix uniformly.
1.4 Early Ideas about the Building Blocks of Matter
Leucippus and his student Democritus proposed matter is composed of small, indestructible particles.
Plato and Aristotle did not embrace atomic ideas; believed matter had no smallest parts and proposed fire, air, earth, and water in varying proportions.
John Dalton offered convincing evidence supporting early atomic ideas. Dalton’s atomic theory of matter includes key postulates:
Each element is composed of tiny, indestructible particles called atoms.
All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements.
Atoms combine in simple, whole-number ratios to form compounds.
Atoms cannot change into atoms of another element in a chemical reaction; reactions involve re-binding of atoms.
1.5 Modern Atomic Theory and the Laws that Led to It
Dalton’s atomic theory explained the following laws:
Each element is composed of tiny, indestructible particles called atoms.
All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements.
Atoms combine in simple, whole-number ratios to form compounds.
In a chemical reaction, atoms cannot be created or destroyed; they are rearranged.
Modern Atomic Theory and Its Laws:
Law of conservation of mass
Law of definite proportions (constant composition)
Law of multiple proportions
The Law of Conservation of Mass
In a chemical reaction, matter is neither created nor destroyed.
When a reaction occurs, the total mass of substances involved does not change.
This is consistent with the idea that matter is composed of small, indestructible particles.
The Law of Definite Proportions (Constant Composition)
All samples of a given compound have the same proportions of constituent elements.
Example: Decomposition of 18.0 g of water (H$2$O) yields 16.0 g O$2$ and 2.0 g H$_2$; O:H mass ratio = 8:1.
The Law of Multiple Proportions
When two elements (A and B) form two different compounds, the masses of B that combine with 1 g of A are in small whole-number ratios.
If A combines with 1, 2, 3, … atoms of B, possible formulas include AB$1$, AB$2$, AB$_3$, etc.
1.6 The Discovery of the Electron
J. J. Thomson’s cathode ray experiments:
Cathode rays traveled from the negatively charged electrode (cathode) to the positively charged electrode (anode).
The particles in the cathode ray:
travel in straight lines
carry a negative electrical charge
Thomson measured the charge-to-mass ratio (e/m) of the cathode ray particles: rac{e}{m} = -1.76 \times 10^{8} \ \text{C/g}.
1.7 The Structure of the Atom
The Early Models:
Thomson’s Plum-Pudding Model: electrons as small negative particles embedded in a positively charged sphere.
Rutherford’s Gold Foil Experiment: positively charged particles directed at a thin gold foil; most pass through, some deflect, a few bounce back.
Conclusion: Matter contains dense regions (nucleus) and large empty space.
Building on Rutherford: The Nuclear Atom Model
Most of the atom’s mass and all of its positive charge are in the nucleus.
Most of the atom’s volume is empty space where negatively charged electrons reside.
The number of electrons outside the nucleus equals the number of protons inside the nucleus; the atom is electrically neutral.
Practice problems involve identifying isotopes, atomic numbers, mass numbers, neutron counts, and electron counts from given symbols.
1.9 Atomic Mass: The Average Mass of an Element’s Atom
Atomic mass is sometimes called atomic weight or standard atomic weight.
The atomic mass listed with an element symbol in the periodic table represents the weighted average mass of the element’s isotopes based on natural abundances.
Formula:
\text{Atomic mass} = \sum_{i} (\text{fraction of isotope } i) \times (\text{mass of isotope } i)