KaiserScience

Some people say not to use ice cubes in soda or wine because “it makes flavor molecules contract”, which supposedly makes a drink taste worse.  Is this correct?

First note that flavor comes from individual molecules dissolved in solution. These are monomers, not polymers. That’s going to be important.

Let’s break the question down into two parts: Does cooling a drink make flavor molecules contract? Does cooling a drink change how our tongue perceives flavor from such molecules?

(A) Does cooling a drink make flavor molecules contract?

Let’s start with the claim that ice makes flavor molecules contract. Sounds reasonable, after all, in everyday life we see that coldness can shrink materials. For instance, a bimetallic strip consists of two different materials. Each has a different expansion coefficient (way it responds to temperature changes.) When heated, one metal expands more than the other, which forces the metal to bend. These strips are used as switches in some thermostats.

image from hyperphysics.phy-astr.gsu.edu

More commonly, we see large scale materials contract when cooled, like highways. That is why roads over bridges, and in parking lots, need thermal expansion joints. Otherwise the shrinking and expansions would otherwise break the surface.

image from Ontario Ministry of Transportation, Bridge Repairs

Characterization of Typical Potent Odorants in Cola-Flavored Carbonated Beverages

Molecules in Coca Cola (Compound Interest)

Molecules in wine (Compound Interest)

Molecules in whiskey (Compound Interest)

But why do these materials contract when cooled? Has each individual molecule shrunk? No. To understand why we need to think about intermolecular vs intramolecular forces.

Intermolecular forces hold one molecule to another.

image from dynamicscience.com.au

Intramolecular forces hold atoms together within a single molecule. (e.g. chemical bonds, covalent bonds)

The naming comes from Latin roots

inter meaning between or among

intra meaning inside

When we use ice to cool a drink from 75 F down to 45 F, does the size of an individual molecule change in any appreciable way? Well, as we see here, a heated molecule can vibrate faster or slower – but the distance between atoms doesn’t appreciably change: the size of any single molecule is constant.

https://en.wikipedia.org/wiki/Thermodynamic_temperature

Nope. First, notice that the temperature in either F or C is misleading. Those temperatures are only measured relative to the freezing point of water. A better comparison is comparing their absolute temperature using the Kelvin scale

75 F = 24 C = 297 K

45 F = 7 C   = 280 K

Wow, on an absolute scale this doesn’t change the temperature of the molecules very much.

Now sure, cooling does reduce molecular vibration: this allows intermolecular forces to pull the molecules closer together, which causes shrinking of the material as a whole. Still, this doesn’t change the bond lengths within a molecule.

(B) Does cooling a drink change how our tongue perceives flavor from such molecules?

If flavor change exists then I would think it is because the tongue tastes certain molecules better at certain temperatures.

image from blogs.unimelb.edu.au, The University of Melbourne, Scientific Scribbles, Let’s talk about Taste

images

https://maxfacts.uk/help/oral-food/ttt

discussion

https://www.ncbi.nlm.nih.gov/books/NBK236241/

https://www.ncbi.nlm.nih.gov/pubmedhealth/PMH0072592/

https://www.livescience.com/20286-foods-taste-hot-cold.html

https://academic.oup.com/chemse/article/42/2/153/2547704

https://www.beveragedaily.com/Article/2005/12/19/Food-temperature-affects-taste-reveal-scientists

https://www.quora.com/How-exactly-does-temperature-affect-the-taste-of-food-and-beverages

Learning Standards

Next Generation Science Standards

HS-PS1-3. Plan and conduct an investigation to gather evidence to compare the structure of substances at the bulk scale to infer the strength of electrical forces between particles. The structure and interactions of matter at the bulk scale are determined by electrical forces within and between atoms.

Massachusetts Science and Technology/Engineering Curriculum Framework

HS-PS1-3. Cite evidence to relate physical properties of substances at the bulk scale to spatial arrangements, movement, and strength of electrostatic forces among ions, small
molecules, or regions of large molecules in the substances. Make arguments to
account for how compositional and structural differences in molecules result in
different types of intermolecular or intramolecular interactions.

Common Core State Standards Connections:

from “An Introduction to Chemistry by Mark Bishop”

http://preparatorychemistry.com/Bishop_pH_Equilibrium.htm

According to the Arrhenius theory of acids and bases, when an acid is added to water, it donates an H⁺ ion to water to form H₃O⁺ (often represented by H⁺).

The higher the concentration of H₃O⁺ (or H⁺) in a solution, the more acidic the solution is.

An Arrhenius base is a substance that generates hydroxide ions, OH^–, in water. The higher the concentration of OH^– in a solution, the more basic the solution is.

Pure water undergoes a reversible reaction in which both H+ and OH^– are generated.

H₂O(l)        H⁺(aq)  +  OH^–(aq)

The equilibrium constant for this reaction, called the water dissociation constant, K_w, is 1.01 × 10^-14 at 25 °C.

K_w  =  [H⁺][OH^–]  =  1.01 × 10^-14   at 25 °C

Because every H⁺ (H₃O⁺) ion that forms is accompanied by the formation of an OH^– ion, the concentrations of these ions in pure water are the same and can be calculated from K_w.

K_w  =  [H⁺][OH^–]  =  (x)(x)  =  1.01 × 10^-14

x  =  [H⁺]  =  [OH-]  =  1.01 × 10^-7 M
(1.005 × 10^-7 M before rounding)

The equilibrium constant expression shows that the concentrations of H⁺ and OH^– in water are linked. As one increases, the other must decrease to keep the product of the concentrations equal to 1.01 × 10^-14 (at 25 °C).

If an acid, like hydrochloric acid, is added to water, the concentration of the H⁺ goes up, and the concentration of the OH^– goes down, but the product of those concentrations remains the same.

An acidic solution can be defined as a solution in which the [H⁺] > [OH^–].

The example below illustrates this relationship between the concentrations of H⁺ and OH^– in an acidic solution.

EXAMPLE 1 – Determining the Molarity of Acids and Bases in Aqueous Solution:  Determine the molarities of H⁺ and OH^– in a 0.025 M HCl solution at 25 °C.

Solution:

K_w  =  [H⁺][OH^–]  =  1.01 × 10^-14   at 25 °C

We assume that hydrochloric acid, HCl(aq), like all strong acids, is completely ionized in water. Thus the concentration of H⁺ is equal to the HCl concentration.

[H⁺] = 0.025 M H⁺

We can calculate the concentration of OH^– by rearranging the water dissociation constant expression to solve for [OH^–] and plugging in 1.01 × 10^-14 for K_w and 0.025 for [H⁺].

Note that the [OH^–] is not zero, even in a dilute acid solution.

If a base, such as sodium hydroxide, is added to water, the concentration of hydroxide goes up, and the concentration of hydronium ion goes down. A basic solution can be defined as a solution in which the [OH^–] > [H⁺].

EXAMPLE 2 – Determining the Molarity of Acids and Bases in Aqueous Solution:    Determine the molarities of H⁺ and OH^– in a 2.9 × 10^-3 M NaOH solution at 30 °C.

Solution:

K_w  =  [H⁺][OH^–]  =  1.47 × 10^-14   at 30 °C    (From Table)

Sodium hydroxide is a water-soluble ionic compound and a strong electrolyte, so we assume that it is completely ionized in water, making the concentration of OH- equal to the NaOH concentration.

[OH^–] = 2.9 × 10^-3 M OH^–

Note that the [H⁺] is not zero even in a dilute solution of base.

Typical solutions of dilute acid or base have concentrations of H⁺ and OH^– between 10^-14 M and 1 M. The table below shows the relationship between the H⁺ and OH^– concentrations in this range.

Concentrations of H⁺ and OH^– in Dilute Acid and Base Solutions at 25 °C

We could describe the relative strengths of dilute solutions of acids and bases by listing the molarity of H⁺ for acidic solutions and the molarity of OH^– for basic solutions. There are two reasons why we use the pH scale instead.

The first reason is that instead of describing acidic solutions with [H⁺] and basic solutions with [OH^–], chemists prefer to have one scale for describing both acidic and basic solutions. Because the product of the H⁺ and OH^– concentrations in such solutions is always 1.01 × 10^-14 at 25 °C, when we give the concentration of H⁺, we are indirectly also giving the concentration of OH^–.

For example, when we say that the concentration of H⁺ in an acidic solution at 25 °C is 10^-3 M, we are indirectly saying that the concentration of OH^– in this same solution is 10^-11 M.

When we say that the concentration of H⁺ in a basic solution at 25 °C is 10^-10 M, we are indirectly saying that the OH^– concentration is 10^-4 M.

The pH concept makes use of this relationship to describe both dilute acid and dilute base solutions on a single scale.

The next reason for using the pH scale instead of H⁺ and OH^– concentrations is that in dilute solutions, the concentration of H⁺ is small, leading to the inconvenience of measurements with many decimal places, such as 0.000001 M H⁺, or to the potential confusion associated with scientific notation, as with 1 × 10^-6 M H⁺.

In order to avoid such inconvenience and possible confusion, pH is defined as the negative logarithm of the H⁺ concentration.

pH  =  -log[H⁺]

Instead of saying that a solution is 0.0000010 M H⁺ (or 1.0 × 10^-6 M H⁺) and 0.000000010 M OH^– (or 1.0 × 10^-8 M OH^–), we can indirectly convey the same information by saying that the pH is 6.00.

pH  =  -log[H⁺]  =  -log(1.0 × 10^-6)  = 6.00

When  taking  the logarithm of a number, report the same number of decimal positions in the answer as you had significant figures in the original value.

Because 1.0 × 10^-6 has two significant figures, we report 6.00 as the pH for a solution with 1.0 × 10^-6 M H⁺.

The table below shows a range of pH values for dilute solutions of acid and base.

pH of Dilute Solutions of Acids and Bases at 25 °C

This table illustrates several important points about pH. Notice that

• When the solution is acidic ([H⁺] > [OH^–), the pH is less than 7.
• When the solution is basic ([OH^–] > [H⁺]), the pH is greater than 7.
• When the solution is neutral ([H⁺] = [OH^–]), the pH is 7. (Solutions with pH’s between 6 and 8 are often considered essentially neutral.)

Also notice that

• As a solution gets more acidic (as [H⁺] increases), the pH decreases.
• As a solution gets more basic (higher [OH^–]), the pH increases.
• As the pH of a solution decreases by one pH unit, the concentration of H⁺ increases by ten times.
• As the pH of a solution increases by one pH unit, the concentration of OH^– increases by ten times.
• The pH, [H⁺], and [OH^–] of some common solutions are listed in the figure below. Notice that gastric juice in our stomach has a pH of about 1.4, and orange juice has a pH of about 2.8. Thus, gastric juice is more than ten times more concentrated in H⁺ than orange juice.

The pH difference of about 4 between household ammonia solutions (pH about 11.9) and milk (pH about 6.9) shows that household ammonia has about ten thousand (10⁴) times the hydroxide concentration of milk.

pH of Common Substances     Acidic solutions have pH values less than 7, and basic solutions have pH values greater than 7. The more acidic the solution is, the lower its pH. The more basic a solution is, the higher the pH.

The corresponding H⁺ and OH^– concentrations are shown in units of molarity. Notice that a decrease of one pH unit corresponds to a ten-fold increase in [H⁺], and an increase of one pH unit for a basic solution corresponds to a ten-fold increase in [OH^–].

EXAMPLE 3 – pH Calculations:  In Example 1, we found that the H⁺ concentration of a 0.025 M HCl solution was 0.025 M H⁺. What is its pH?

Solution:

pH = -log[H⁺]  =  -log(0.025) =  1.60

EXAMPLE 4 – pH Calculations:  In Example 2, we found that the H⁺ concentration of a 2.9 × 10^-3 NaOH solution was 5.1 × 10^-12 M H⁺. What is its pH?

Solution: 

pH = -log[H⁺]  =  -log(7.5 × 10^-12) =  11.29

We can convert from pH to [H⁺] and [OH^–] using the following equations, as demonstrated in Examples 5 and 6.

[H⁺]  =  10^-pH

EXAMPLE 5 – pH Calculations:  What is the [H⁺] in a glass of lemon juice with a pH of 2.12?

Solution:

[H⁺]  =  10^-pH  =  10^-2.12  =  7.6 × 10^-3 M H⁺

EXAMPLE 6  – pH Calculations: What is the [OH^–] in a container of household ammonia at 25 °C with a pH of 11.900?

Solution:

[H⁺]  =  10^-pH  =  10-11.900  =  1.26 × 10^-12 M H⁺

Learning Standards

2016 Massachusetts Science and Technology/Engineering Curriculum Framework

HS-PS1-9 (MA). Relate the strength of an aqueous acidic or basic solution to the extent of an acid or base reacting with water, as measured by the hydronium ion concentration (pH) of the solution. Make arguments about the relative strengths of two acids or bases with similar structure and composition.

Science and Engineering Practices

Mathematical and computational thinking in 9–12 builds on pre-K–8 and experiences and progresses to using algebraic thinking and analysis, a range of linear and nonlinear functions including trigonometric functions, exponentials and logarithms, and computational tools for statistical analysis to analyze, represent, and model data.

Benchmarks for Science Literacy, AAAS

Most cells function best within a narrow range of temperature and acidity. At very low temperatures, reaction rates are too slow. High temperatures and/or extremes of acidity can irreversibly change the structure of most protein molecules. Even small changes in acidity can alter the molecules and how they interact. 5C/H7

The temperature and acidity of a solution influence reaction rates. Many substances dissolve in water, which may greatly facilitate reactions between them. 4D/M4

ACS Middle School Chemistry Lessons

From middleschoolchemistry.com, contact staff at ACS. Copyright 2015 American Chemical Society

Online textbook: Chapter 5: Acids Bases and their reactions

http://www.chem4kids.com/files/react_acidbase.html

http://www.chemistryland.com/CHM151S/04-Solutions/acids/AcidsBases151.html